Higher Chemistry - Unit 1: Chemical Changes and Structure

1.1 Periodicity PPQ

Table of Contents
  1. SQA Specifications
  2. N5 PPQ Revision
  3. The First 20 Elements
  4. Covalent Radius
  5. Ionisation Energy
  6. Electronegativity
SQA Specifications
  • Elements are arranged in the periodic table in order of increasing atomic number. The periodic table is not just an organization of elements but a tool that enables chemists to make accurate predictions concerning the physical properties and chemical behaviour of any element based on its position within the table.
Key Features of the Periodic Table:
  • Groups: Vertical columns; elements within the same group share similar chemical properties due to having a common number of electrons in their outer shell.
  • Periods: Horizontal rows; elements are organized with increasing atomic number which signifies an escalating number of outer electrons and a transition from metallic to non-metallic characteristics.
The First 20 Elements
  • The first 20 elements in the periodic table are categorized according to their bonding and structure:
    • Metallic: Lithium (Li), Beryllium (Be), Sodium (Na), Magnesium (Mg), Aluminium (Al), Potassium (K), Calcium (Ca).
    • Covalent Molecular: Hydrogen (H₂), Nitrogen (N₂), Oxygen (O₂), Fluorine (F₂), Chlorine (Cl₂), Phosphorus (P₄), Sulfur (S₈), and Fullerenes (e.g., C₆₀).
    • Covalent Network: Boron (B), Carbon (C) in forms like diamond and graphite, Silicon (Si).
    • Monatomic: Noble gases.
Covalent Radius
  • The covalent radius is defined as a measure of the size of an atom. Trends in covalent radius across periods and down groups can be explained in terms of the number of occupied electron shells and the effective nuclear charge.
Ionisation Energy
  • Ionisation Energy: The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms. Subsequent ionisation energies refer to the removal of further moles of electrons. The trends in ionisation energies moving across periods and down groups can be attributed to changes in atomic size, nuclear charge, and the screening effect introduced by inner-shell electrons.
Electronegativity
  • Electronegativity: Atoms of different elements exhibit varying attractions for bonding electrons. Electronegativity quantitatively measures an atom's ability to attract bonding electrons within a chemical bond. Trends in electronegativity across periods and down groups can similarly be rationalized based on covalent radius, nuclear charge, and the screening effect from inner-shell electrons.

N5 PPQ Revision

  1. Which of the following elements usually exists as diatomic molecules?
    A. Helium
    B. Nitrogen
    C. Silicon
    D. Sulfur

  2. Tennessine is a newly discovered element with a predicted electron arrangement of 2,8,18,32,32,18,7. In which group of the periodic table should Tennessine be placed?
    A. 1
    B. 2
    C. 7
    D. 8

  3. Identify the particle which is a negative ion from the given table of particles.

  4. Identify the positively charged ion from the selections provided.

  5. In which of the following compounds do both ions have the same electron arrangement?
    A. Na₂O
    B. LiF
    C. KBr
    D. MgCl₂

  6. Identify the correct location of a proton and an electron in an atom from the given options.

  7. Which of the following elements does not exist as diatomic molecules?
    A. Oxygen
    B. Helium
    C. Bromine
    D. Hydrogen

  8. From the list below, which element forms an ion with a single positive charge and an electron arrangement of 2,8?
    A. Sodium
    B. Magnesium
    C. Fluorine
    D. Neon

  9. An atom has 21 protons, 21 electrons, and 24 neutrons. The atom has:
    A. Atomic number 24 and mass number 42
    B. Atomic number 45 and mass number 21
    C. Atomic number 21 and mass number 45
    D. Atomic number 24 and mass number 45

  10. When chlorine reacts with sodium, the ionic compound sodium chloride is formed. A chloride ion has a stable electron arrangement. Describe how a chlorine atom achieves this stable electron arrangement.

  11. State what is meant by the term isotope.

  12. Complete the table for Br₃₅₇.

  13. A sample of bromine consists of two isotopes, Br₃₅₇ and Br₃₅₉. The average atomic mass of the sample is 80. This indicates a ratio of the two isotopes in the sample. Explain this.

  14. A sample of tin contains three different isotopes. The nuclide notation for each is shown. Another isotope of tin exists with 74 neutrons. Write the nuclide notation for this isotope of tin.

  15. A molecule of phosphorus trifluoride is shown. Identify its shape from the options:
    A. Linear
    B. Angular
    C. Tetrahedral
    D. Trigonal pyramidal

  16. The shape of a molecule of hydrogen bromide is likely to be:
    A. Tetrahedral
    B. Trigonal pyramidal
    C. Angular
    D. Linear

  17. Which of the following molecules has a trigonal pyramidal shape?
    A. HCl
    B. CO₂
    C. NCl₃
    D. CHCl₃

  18. State the name used to describe the shape of a molecule of chloromethane.

  19. Ionic compounds conduct electricity when molten due to:
    A. Ions that are free to move
    B. Delocalised electrons
    C. Metal atoms
    D. A lattice structure

  20. Covalent and ionic compounds exhibit different physical properties. Complete the table by circling the appropriate descriptors for each type.

  21. There are many different types of glass, primarily composed of silica (SiO₂) derived from sand. Silica possesses a melting point of 1713 °C. State the structural term used to describe silica.

  22. Tin(IV) chloride can be synthesized by reacting tin with chlorine. Some properties of tin(IV) chloride are detailed in the accompanying table. Utilize this information to ascertain the type of bonding present in tin(IV) chloride.

The First 20 Elements

  1. An element with covalent bonding and London dispersion forces. Possible candidates include:
    (1) A. Boron
    (B). Neon
    (C). Sodium
    (D). Sulfur

  2. Identify which structures are never found in compounds:
    (1) A. Covalent molecular
    (B). Covalent network
    (C). Monoatomic
    (D). Ionic

  3. Complete the bonding and structural information for the following elements in the third period of the Periodic Table:

    ElementAluminiumSiliconPhosphorusSulfur
    BondingCovalentCovalent
    StructureLatticeMolecular

  4. Identify the element from the third period that exists as a covalent network.

  5. Explain the increasing boiling points of halogens as they descend the group. In your response, identify the intermolecular forces involved.

Covalent Radius

  1. The dimension in size between sodium and chlorine atoms is predominantly attributed to:
    (1) A. Mass of each atom
    (B). Number of electrons
    (C). Number of neutrons
    (D). Number of protons

  2. Determine which statement is accurate:
    (1) A. The sodium atom is larger than the sodium ion.
    (B). The chloride ion is smaller than the chlorine atom.
    (C). The magnesium ion is larger than the magnesium atom.
    (D). The oxygen atom is larger than the oxide ion.

  3. Identify the expected line in the data table for the element Francium.

  4. Explain the decrease in atomic size across the period from Aluminium to Sulfur. (1)

  5. Detail why the covalent radius of sulfur is smaller than that of phosphorus. (1)

  6. Explain fully why the covalent radius of sodium exceeds that of the sodium ion. (2)

  7. Discuss the reduction of atom size progressing across the third period from sodium to argon. (1)

  8. The ionic radius represents the dimension of an ion. Discuss why the ionic radius of phosphorus is greater compared to that of aluminium. (2)

  9. Nitrogen, oxygen, and fluorine are observed in the second period of the periodic table. Discuss the rationale for the decreasing covalent radius from nitrogen to fluorine.

Ionisation Energy

  1. A graph depicting successive ionisation energies for element Z shows a spike. Identify the group in which element Z exists:
    A. 1
    B. 3
    C. 4
    D. 6

  2. Which formula represents the first ionisation energy of fluorine?
    (1)
    A. F⁻(g) → F(g) + e⁻
    B. F⁻(g) → 1/2F₂(g) + e⁻
    C. F(g) → F⁺(g) + e⁻
    D. 1/2F₂(g) → F⁺(g) + e⁻

  3. The first three ionisation energies of aluminium are presented in a table. Using this data, ascertain the enthalpy change, expressed in kJ mol⁻₁, for the reaction:
    Al⁺(g) → Al³⁺(g) + 2e⁻
    A. 1817
    B. 2395
    C. 4562
    D. 5140

  4. Determine the enthalpy change in kJ mol⁻₁ for the reaction:
    Be(g) → Be²⁺(g) + 2e⁻
    A. 900
    B. 1757
    C. 2657
    D. 3514

  5. From the graph representing energy equivalents of ionisation energies, deduce the most probable ion formed having a charge of:
    A. 2⁺
    B. 3⁺
    C. 2⁻
    D. 3⁻

  6. The difference between the first ionisation energies of sodium and chlorine is primarily due to:
    A. Number of electrons
    B. Number of neutrons
    C. Number of protons
    D. Mass of each atom

  7. Discuss why the first ionisation energy decreases as one moves down Group 1. (1)

  8. Justify why the second ionisation energy significantly exceeds the first for Group 1 elements. (2)

  9. The variation of ionisation energy across a period is notable. Discuss why the first ionisation energy increases across the period. (1)

  10. Write the equation for the second ionisation energy of magnesium with appropriate state symbols.

  11. The tabulated first four ionisation energies of aluminium reveal a significant gap between the third and fourth ionisation energies. Elucidate why this occurs. (1)

  12. Explain why the first ionisation energy escalates from lithium to neon. (1)

  13. Contrast the first ionisation energy of potassium with that of lithium, explaining why potassium’s is lower. (1)

  14. Write the equation representing the first ionisation energy of phosphorus. (1)

  15. Discuss the substantial increase noted between the fifth and sixth ionisation energies of nitrogen. (2)

  16. Define the term first ionisation energy. (1)

  17. Describe the trend of first ionisation energy for group 7 elements as you descend the group. (1)

  18. A diagram showcases a graph of the first ionisation energy against atomic number for several periodic elements. Explain the observed increase in first ionisation energy from elements d to k in the diagram. (1)

  19. From elements a to m in the diagram, state an example of a group 7 element. (1)

  20. The table presents four ionisation energies of sodium. Elaborate on why there is a significant jump between the first and second ionisation energies. (1)

  21. Using the previous table data, calculate the enthalpy change in kJ mol⁻¹ for the following reaction:
    Na⁺(g) ⟶ Na³⁺(g) + 2e⁻

  22. Write an equation indicating the second ionisation energy of nitrogen.

Electronegativity

  1. Identify which element has the greatest attraction for bonding electrons from the following options:
    A. Bromine
    B. Chlorine
    C. Lithium
    D. Sodium

  2. For elements in group 7 of the periodic table, ascertain which statement is true as one descends the group:
    A. The boiling point decreases.
    B. The covalent radius decreases.
    C. Electronegativity decreases.
    D. The strength of London dispersion forces decreases.

  3. Identify which of the following atoms has the least attraction for bonding electrons:
    A. Carbon
    B. Nitrogen
    C. Phosphorus
    D. Silicon

  4. Of the given options, which atom exhibits the greatest attraction for bonding electrons?
    A. Sulfur
    B. Silicon
    C. Nitrogen
    D. Hydrogen

  5. Define the term electronegativity.

  6. State the trend in electronegativity as you move from lithium to fluorine across period 2.

  7. Explain why electronegativity values are absent for elements with atomic numbers 2, 10, and 18.

  8. Electronegativity values play a crucial role in predicting the type of bonding present in substances. Using the accompanying diagram, identify the highest average electronegativity found in ionic compounds.

  9. Elaborate on why electronegativity diminishes as one descends a group. (2)

  10. Suggest which of the group 2 elements serves as the most effective reducing agent.