Misconceptions

The following is a list of common misconceptions that students may have around the topics in this unit. Included with the misconceptions are the actual facts, ways to overcome the misconception and some activity or lesson ideas that can be used.

Figure B.3: Misconception Cartoon

1. The elements on the periodic table are arranged by increasing atomic mass

  • How to overcome this misconception
    • It may seem like the elements are arranged by increasing atomic number but the elements are arranged by increasing atomic number.
    • Atomic mass is based on the mass of the entire atom and can vary due to a number of circumstances. Atomic number is the number or protons which is unique to each atom
  • Activity/Lesson Ideas
    • Have students perform a “scavenger hunt” using the periodic table to find all of the instances where the elements are not arranged by increasing atomic mass

2. Low ionization energy means that it is hard to remove that electron.

  • How to overcome this misconception
    • Explain to students that ionization energy is the amount of energy required to remove an electron from that atom or molecule.
    • The lower the energy required to remove an electron, the easier it is to remove.
  • Activity/Lesson Ideas
    • Bring in items that are stuck together (e.g., two lego blocks, two magnets, two items super-glued together, two items taped together)
    • Have students draw a parallel between the amount of energy they had to put in to pull those two items apart and how easy or hard it was to do so.
  • https://www.ionicviper.org/class-activity/athletic-periodic-trends-review

3. Atomic radius increases with increasing atomic number and mass

  • How to overcome this misconception
    • This misconception occurs when students do not understand the reasoning behind the trend
    • Atomic radius increases going down each family because the attractive force of the nucleus is reduced due to shielding. This allows for a larger radius because the electrons are less attracted to the nucleus.
    • Atomic radius decreases going from left to right along each period. Nuclear charge increases as you move along, however the amount of shielding remains constant. Therefore the strength of the attraction between the nucleus and valence electrons is stronger, decreasing the size of the atom.
  • Activity/Lesson Ideas
    • Have students determine an explanation for the fact that Nitrogen’s atomic radius is smaller than both Carbon and Oxygen.
  • http://faculty.bmcc.cuny.edu/faculty/upload/Article-atomic-properties.pdf

4. The most current model of the atom replaces all previous models.

  • The Schrödinger atomic model does not replace all previous models. It describes the location of the electrons orbiting the nucleus more accurately than other models. However, the Bohr model is more useful when predicting chemical bonding and describing chemical properties of elements.
  • Douglass, S. et al. (2010). Chemistry 11: Teacher resource. McGraw-Hill Ryerson.

5. The first 92 elements in the periodic table have naturally occurring stable isotopes.

  • This is a general statement, but there are two elements that do not have stable isotopes - 43 (Tc) and 61 (Pm).
  • Activity/Lesson Ideas
    • Have students research which elements do not have stable isotopes and share this information with their classmates.
  • Douglass, S. et al. (2010). Chemistry 11: Teacher resource. McGraw-Hill Ryerson.

6. The periodic table was constructed using the atomic structure of the elements.

  • Mendeleev first constructed the periodic table using empirical evidence about atomic mass and properties of elements. In 1914 Moseley developed a method to determine the number of protons in an element. This allowed scientists to relate atomic structure to the organization proposed by Mendeleev.
  • Activity/Lesson Ideas
  • Douglass, S. et al. (2010). Chemistry 11: Teacher resource. McGraw-Hill Ryerson.

7. The electron affinity is the same quantity as the ionization energy.

  • Both quantities occur in relationship to a neutral atom. The amount of energy that is released when an atom gains an electron would be the same as the amount of energy required to remove that exact electron from the ion, not from a neutral atom.
  • Activity/Lesson Ideas
    • Have students read about the electron affinity and ionization energy of a specific element. Students will see that the quantities are different and then the teacher can have them try to explain why.
  • Douglass, S. et al. (2010). Chemistry 11: Teacher resource. McGraw-Hill Ryerson.

8. Elements with the greatest electron affinity have the highest electronegativity.

  • This correlation is not absolutely consistent. For example, Flourine, which has the highest electronegativity releases less energy when it gains an electron compared to chlorine, which has a lower electronegativity.
  • Douglass, S. et al. (2010). Chemistry 11: Teacher resource. McGraw-Hill Ryerson.

9. The term salt refers only to table salt, NaCl.

  • The term salt refers to all ionic compounds.
  • Douglass, S. et al. (2010). Chemistry 11: Teacher resource. McGraw-Hill Ryerson.

10. Ionic bonds are formed when electron transfer from one atom to to another.

  • Metal atoms lose electrons and become positively charged while non-metal atoms gain electrons and become negatively charged. These oppositely charged ions are electrostatically attracted to each other, this results in an ionic bond.
  • Activity/Lesson Ideas
    • To overcome this you can draw simple diagrams on the chalkboard, whiteboard or interactive whiteboard to illustrate the formation of the ionic bond.
  • Douglass, S. et al. (2010). Chemistry 11: Teacher resource. McGraw-Hill Ryerson.

Figure B.4: Ionic Bonding

11. Ionic bonds form between metals and non-metals, while covalent bonds form between non-metals.

  • The true nature of a chemical bond is determined using the differences in electronegativities. Bonds have "ionic character" or "covalent character" along a bonding continuum.
  • Douglass, S. et al. (2010). Chemistry 11: Teacher resource. McGraw-Hill Ryerson.

Figure B.5: Bonding Continuum

12. When a substance melts or boils, bonds are broken.

  • Bonds are only broken when an ionic compound melts or boils. In a molecular compound, the covalent bonds remain undamaged, only the forces holding the molecules together with other molecules are broken.
  • Douglass, S. et al. (2010). Chemistry 11: Teacher resource. McGraw-Hill Ryerson.
  • Activity/Lesson Ideas
    • Review the states of matter with students and explain that bonds are not broken during a change of state.
    • The following video might be useful:

13. All ionic compounds conduct an electric current.

  • When ionic compounds are in the solid state, the ions are held rigidly together and do not conduct electricity. When ionic compounds are in the liquid state, they will conduct electricity.
  • Douglass, S. et al. (2010). Chemistry 11: Teacher resource. McGraw-Hill Ryerson.
  • Activity/Lesson Ideas
    • Have students use a conductivity apparatus and try to test if solid table salt (NaCl) conducts electricity. They will notice that it does not. Next, have students dissolve table salt in water and use the conductivity apparatus to test if aqueous NaCl conducts electricity.

Figure B.6: Salt Figure B.7: Dissolved Salt

Next: Lesson Sequence