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Structure 1.1 Flashcards

Models of the Particulate Nature of Matter. Select a sub-topic below

1.1.1 Elements, Compounds & Mixtures: read all 8 questions and answers
  1. Define an element.
    A substance that cannot be chemically broken down into simpler substances. All atoms have the same atomic number (same number of protons).
  2. Define a compound.
    A substance formed when atoms of two or more different elements are chemically bonded together in a fixed, whole-number ratio.
  3. State the law of constant composition.
    A given compound always contains the same elements combined in the same proportions by mass, regardless of its source or method of preparation.
  4. What is a mixture?
    Two or more substances (elements or compounds) that are not chemically bonded. They retain their own properties and can be separated by physical methods.
  5. What is a homogeneous mixture?
    A mixture with uniform composition throughout, appearing as a single phase. Examples: salt water, air, alloys.
  6. What is a heterogeneous mixture?
    A mixture with non-uniform composition and visible phase boundaries. Examples: sand and water, oil and water.
  7. Why are alloys classified as mixtures, not compounds?
    Alloys contain metallic elements in a variable, non-fixed ratio determined by manufacturing, not by stoichiometric laws.
  8. Name 5 physical separation techniques.
    Filtration, evaporation, distillation, recrystallisation, chromatography.
1.1.2 Kinetic Molecular Theory: read all 6 questions and answers
  1. Describe particle arrangement in a solid.
    Particles are closely packed in a regular arrangement. They vibrate around fixed positions but cannot move freely.
  2. Describe particle arrangement in a liquid.
    Particles are close together but in an irregular arrangement. They can slide past each other and move around.
  3. Describe particle arrangement in a gas.
    Particles are far apart with no regular arrangement. They move randomly at high speed in all directions.
  4. What happens to particles during melting?
    Particles gain enough energy to overcome some intermolecular forces. They break free from fixed positions but remain close together. Temperature stays constant during the phase change.
  5. What happens to particles during boiling?
    Particles gain enough energy to completely overcome intermolecular forces. They escape the liquid phase and move freely as gas. Temperature stays constant during the phase change.
  6. Why does temperature remain constant during a phase change?
    All energy input is used to overcome intermolecular forces (potential energy increases) rather than increasing kinetic energy of the particles.
1.1.3 Temperature & Kelvin: read all 4 questions and answers
  1. Convert Celsius to Kelvin.
    K = degrees C + 273.15 (in IB, rounding to 273 is often acceptable).
  2. What is absolute zero?
    0 K (or -273.15 degrees C). The lowest possible temperature where particles have minimum kinetic energy and (theoretically) cease all motion.
  3. Why do we use the Kelvin scale in chemistry?
    Kelvin is directly proportional to kinetic energy of particles. Doubling the Kelvin temperature doubles the average kinetic energy. The Celsius scale does not have this proportionality.
  4. What is the relationship between temperature and kinetic energy?
    Average kinetic energy of particles is directly proportional to the absolute (Kelvin) temperature: Ek is proportional to T (in Kelvin).
1.2.1 Subatomic Particles: read all 6 questions and answers
  1. What are the three subatomic particles?
    Proton (charge +1, mass 1 amu), Neutron (charge 0, mass 1 amu), Electron (charge -1, mass ~1/1836 amu).
  2. What is the Atomic Number (Z)?
    The number of protons in the nucleus. It defines the element and determines its position in the periodic table.
  3. What is the Mass Number (A)?
    The total number of protons and neutrons in the nucleus. A = Z + number of neutrons.
  4. Where is virtually all the mass of an atom located?
    In the nucleus, which contains protons and neutrons. The electrons contribute negligible mass.
  5. How is the number of neutrons calculated?
    Number of neutrons = Mass Number (A) - Atomic Number (Z).
  6. What determines the chemical properties of an element?
    The electron configuration, which is determined by the number of protons (atomic number).
1.2.2 Isotopes & Mass Spectrometry: read all 6 questions and answers
  1. Define isotopes.
    Atoms of the same element (same number of protons/atomic number) with different numbers of neutrons (different mass number).
  2. Why do isotopes have identical chemical properties?
    Because they have the same electron configuration. Chemical properties depend on electrons, not neutrons.
  3. Why do isotopes have different physical properties?
    Because they have different masses (different numbers of neutrons), affecting properties like density, rate of diffusion, and melting/boiling points.
  4. List the 4 stages of mass spectrometry.
    1. Vaporisation, 2. Ionisation (electron bombardment), 3. Acceleration (electric field), 4. Deflection (magnetic field), then Detection.
  5. How is relative atomic mass (Ar) calculated from mass spec data?
    Ar = sum of (isotope mass x percentage abundance) / 100.
  6. What does the height of each peak in a mass spectrum represent?
    The relative abundance of that isotope.
1.2.3 Emission Spectra: read all 6 questions and answers
  1. What causes an atomic emission spectrum?
    Electrons absorb energy and are excited to higher energy levels. When they fall back down, they emit photons of specific frequencies, producing a line spectrum.
  2. Why do emission spectra have discrete lines?
    Because electron energy levels are quantised. Only specific energy transitions are possible, producing photons of specific frequencies.
  3. What does convergence of spectral lines indicate?
    It indicates that energy levels get closer together at higher energies. The convergence limit corresponds to ionisation energy.
  4. How can the first ionisation energy be determined from an emission spectrum?
    From the convergence limit of the highest-energy series (Lyman series for hydrogen). The frequency at convergence corresponds to the ionisation energy.
  5. What is the difference between a continuous and a line spectrum?
    A continuous spectrum contains all wavelengths (e.g. White light). A line spectrum contains only specific wavelengths, produced by excited atoms.
  6. Why is the hydrogen emission spectrum important?
    It provides evidence for quantised energy levels and was key to developing the Bohr model of the atom.
1.3.1 Shells, Subshells & Orbitals: read all 6 questions and answers
  1. What is the maximum number of electrons in each main energy level (n)?
    2n squared. So: n=1 holds 2, n=2 holds 8, n=3 holds 18, n=4 holds 32.
  2. What are the four types of subshell?
    s (holds 2e), p (holds 6e), d (holds 10e), f (holds 14e).
  3. What is an orbital?
    A region of space around the nucleus where there is a high probability of finding an electron. Each orbital holds a maximum of 2 electrons with opposite spins.
  4. How many orbitals in each subshell type?
    s: 1 orbital, p: 3 orbitals, d: 5 orbitals, f: 7 orbitals.
  5. Describe the shape of an s orbital.
    Spherical, centred on the nucleus.
  6. Describe the shape of a p orbital.
    Dumbbell-shaped (two lobes). The three p orbitals are oriented along x, y, and z axes at 90 degrees to each other.
1.3.2 Aufbau, Hund & Pauli: read all 6 questions and answers
  1. State the Aufbau Principle.
    Electrons fill the lowest available energy orbital first. Order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p...
  2. State Hunds Rule.
    Orbitals of equal energy (degenerate) are each singly occupied before any orbital is doubly occupied, and all singly occupied orbitals have the same spin.
  3. State the Pauli Exclusion Principle.
    No two electrons in an atom can have the same set of four quantum numbers. In practice, a maximum of 2 electrons per orbital with opposite spins.
  4. Why is the electron configuration of Chromium anomalous?
    [Ar] 3d5 4s1 (not [Ar] 3d4 4s2). A half-filled d subshell provides extra stability due to symmetrical electron distribution.
  5. Why is the electron configuration of Copper anomalous?
    [Ar] 3d10 4s1 (not [Ar] 3d9 4s2). A fully filled d subshell provides extra stability.
  6. Why does 4s fill before 3d?
    In neutral atoms, the 4s orbital has slightly lower energy than 3d. However, once the atom is ionised, 3d becomes lower in energy than 4s.
1.3.3 Ion Configurations: read all 5 questions and answers
  1. Which electrons are removed first when forming transition metal cations?
    The 4s electrons are removed before 3d electrons. E.g. Fe2+ is [Ar] 3d6 (not [Ar] 3d4 4s2).
  2. Write the electron configuration of Fe2+.
    [Ar] 3d6. Iron loses its two 4s electrons first.
  3. Write the electron configuration of Fe3+.
    [Ar] 3d5. Iron loses two 4s electrons and one 3d electron.
  4. How do negative ions form?
    Atoms gain electrons to achieve a full outer shell. E.g. O2- is 1s2 2s2 2p6 (same as Ne).
  5. What are isoelectronic species?
    Species with the same electron configuration. E.g. Na+, Mg2+, F-, O2-, and Ne all have the configuration 1s2 2s2 2p6.
1.3.4 Ionisation Energy Trends: read all 6 questions and answers
  1. Define first ionisation energy.
    The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
  2. How does IE trend across a period?
    Generally increases across a period due to increasing nuclear charge with the same shielding, so greater attraction for outer electrons.
  3. How does IE trend down a group?
    Decreases down a group because outer electrons are further from the nucleus with more shielding, so less attraction.
  4. Why does IE drop between Groups 2 and 13?
    Group 13 elements lose a p electron, which is higher in energy and easier to remove than the s electron lost by Group 2.
  5. Why does IE drop between Groups 15 and 16?
    In Group 16, the electron being removed is paired in a p orbital. Electron-electron repulsion in the pair makes it easier to remove.
  6. How do successive ionisation energies provide evidence for electron shells?
    Large jumps in IE occur when electrons are removed from a new (inner) shell, indicating distinct energy levels.
1.4.1 The Mole Concept: read all 6 questions and answers
  1. What is Avogadro's Constant?
    6.02 × 10²³ particles per mole (in the IB data booklet; defined in SI as exactly 6.02214076 × 10²³ mol⁻¹). Historically, it was the number of atoms in exactly 12 g of carbon-12.
  2. State the formula for moles from mass.
    n = m / M, where n = moles, m = mass in grams, M = molar mass in g/mol.
  3. What is molar mass?
    The mass of one mole of a substance, in g/mol. Numerically equal to the relative molecular/formula mass.
  4. What is the molar volume of a gas at STP?
    22.7 dm³ mol⁻¹ (at 273.15 K and 100 kPa). This applies to all ideal gases.
  5. State the formula for moles of gas at STP.
    n = V / 22.7, where V is volume in dm³ at STP.
  6. What is the formula for number of particles?
    Number of particles = \(n \times 6.02 \times 10^{23}\) (moles times Avogadro's constant).
1.4.2 Empirical & Molecular Formulas: read all 6 questions and answers
  1. What is an empirical formula?
    The simplest whole-number ratio of atoms of each element in a compound. E.g. CH2O for glucose.
  2. What is a molecular formula?
    The actual number of atoms of each element in one molecule. E.g. C6H12O6 for glucose.
  3. How do you find the molecular formula from the empirical formula?
    Divide the molar mass of the compound by the molar mass of the empirical formula. Multiply the empirical formula by this whole number.
  4. Define percentage composition by mass.
    The mass of each element in a compound expressed as a percentage of the total molar mass.
  5. What is a limiting reagent?
    The reactant that is completely consumed first in a reaction, determining the maximum amount of product that can form.
  6. How do you identify the limiting reagent?
    Calculate moles of each reactant, divide by the stoichiometric coefficient, and the smallest value identifies the limiting reagent.
1.4.3 Solutions & Concentration: read all 6 questions and answers
  1. State the formula for concentration.
    c = n / V, where c = concentration in mol dm⁻³, n = moles, V = volume in dm³.
  2. How do you convert cm³ to dm³?
    Divide by 1000. So 250 cm³ = 0.250 dm³.
  3. What is a standard solution?
    A solution of known, precisely determined concentration, prepared by dissolving an accurate mass of solute and making up to an exact volume of solution in a volumetric flask.
  4. What is the dilution formula?
    c1V1 = c2V2, where c1 and V1 are initial concentration and volume, c2 and V2 are final.
  5. What does stoichiometric coefficient mean?
    The number in front of a formula in a balanced equation, showing the molar ratio of reactants and products.
  6. What is percentage yield?
    Percentage yield = (actual yield / theoretical yield) x 100%. In practice it is often less than 100% due to side reactions, mechanical transfer losses, or incomplete reactions (apparent yields > 100% usually indicate wet or impure product).
1.5.1 Gas Laws: read all 6 questions and answers
  1. State Boyles Law.
    At constant temperature, the pressure of a fixed mass of gas is inversely proportional to its volume. P1V1 = P2V2.
  2. State Charless Law.
    At constant pressure, the volume of a fixed mass of gas is directly proportional to its absolute temperature (in Kelvin). V1/T1 = V2/T2.
  3. State Gay-Lussacs Law.
    At constant volume, the pressure of a fixed mass of gas is directly proportional to its absolute temperature (in Kelvin). P1/T1 = P2/T2.
  4. State Avogadros Law.
    Equal volumes of gases at the same temperature and pressure contain equal numbers of particles (moles). V is proportional to n.
  5. What is STP in the IB?
    Standard Temperature and Pressure: 273.15 K (0 degrees C) and 100 kPa.
  6. What is the combined gas law?
    P1V1/T1 = P2V2/T2 (combines Boyles, Charless, and Gay-Lussacs laws for a fixed amount of gas).
1.5.2 Ideal Gas Equation: read all 6 questions and answers
  1. State the ideal gas equation.
    PV = nRT, where P = pressure (Pa), V = volume (m³), n = moles, \(R = 8.314\text{ J mol}^{-1}\text{ K}^{-1}\), T = temperature (K).
  2. What are the SI units for PV = nRT?
    P in Pa (not kPa), V in m³ (not dm³ or cm³), T in K (not degrees C), \(R = 8.314\text{ J mol}^{-1}\text{ K}^{-1}\).
  3. How do you convert kPa to Pa?
    Multiply by 1000. So 101.3 kPa = 101300 Pa.
  4. How do you convert dm³ to m³?
    Divide by 1000. So 22.7 dm³ = 0.0227 m³.
  5. How can you use PV = nRT to find molar mass?
    Since n = m/M, rearrange to M = mRT/(PV). Measure mass, temperature, pressure, and volume to find M.
  6. What is the value of the gas constant R?
    \(R = 8.314\text{ J mol}^{-1}\text{ K}^{-1}\). This is given in the IB data booklet.
1.5.3 Real vs Ideal Gases: read all 6 questions and answers
  1. What are the two key assumptions of the ideal gas model?
    1. Gas particles have negligible volume compared to the container. 2. There are no intermolecular forces between particles.
  2. Under what conditions do real gases behave most ideally?
    At high temperature (particles have high kinetic energy to overcome IMFs) and low pressure (particles are far apart, so volume and IMFs are negligible).
  3. Why do real gases deviate from ideal behaviour at high pressure?
    At high pressure, particles are forced close together. Their actual volume becomes significant compared to container volume, and IMFs become stronger.
  4. Why do real gases deviate from ideal behaviour at low temperature?
    At low temperature, particles move slowly and have less kinetic energy. IMFs become significant relative to kinetic energy, reducing the effective pressure.
  5. Which type of gas behaves most ideally?
    Small, non-polar gases (e.g. He, H2) because they have the weakest IMFs and smallest volume. Large polar molecules deviate the most.
  6. What happens to all real gases at sufficiently low temperature?
    They liquefy (condense), which an ideal gas would never do. This is because IMFs become dominant over kinetic energy.
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