Week Three Wrap-Up
Module Review
This week, we took a closer look at the periodic table and all the information that it contains about elements. Within this table, elements are arranged based on their behaviors and reactions with other elements and are grouped into families: Alkali Metals (1A), Alkaline Earth Metals (2A), Noble Gases, and Halogens. We discovered that elements (for representative elements only) within a family have similar electron configurations which cause them to behave similarly. We learned that the group (family) numbers equal the number of outer electrons (valence electrons), which are the electrons involved in chemical bonding. The rows of the Periodic Table are called Periods and the Shell of an element that the valence electrons are in is equal to the Period number of that element. Wow, this is a lot!
After learning how atoms are arranged in the periodic table of elements, we worked out how to write the electron configuration of anions and cations. We found that an anion is a negatively charged ion and is usually formed from a nonmetal, the negative charge of which gives the atom a noble gas configuration. The electron configuration of anions are written for the atom with the additional electrons placed in the next available orbital. A rule for writing cation (positively charged ions) electronic configurations is always start with the neutral element and then take the electrons from the subshell with the highest n (p before s).
We then learned about effective nuclear charge and the role it plays in the size of the atoms, specifically ions. We talked about how shielding (the effect in which electrons can reduce the electrostatic attraction between the protons of the nucleus and other electrons) reduces the nuclear charge felt by an electron (effective nuclear charge). We used the various strengths of shielding, the strongest being electrons blocked by other electrons in a lower subshell (closer to the nucleus), the slightly strong being electrons competing with other electrons within the same shell, and the lack of shielding by outer electrons (higher shells), to help define the atomic and, ultimately, the ionic radius of atoms. In other words, the size of an atom is ultimately defined by the attraction of the electrons to the protons in the nucleus of the atom; the more the electrons in the outer shells are feeling the nuclear charge (the higher the effective nuclear charge), the smaller the atom will be.
We then defined ionization energy (endothermic reaction) and looked at the role of effective nuclear energy in terms of the trends of ionization energy as we go across the periodic table or up and down within a family within the periodic table. We learned that ionization energy, always a positive value, is the minimum energy required to remove an electron from a gaseous atom in its grounds state and is the measure of how tightly the electrons are held in the atom. As we continue removing electrons from an atom, the energy required will increase. Also, as we move across a period, or up a family, within the periodic table, the ionization energy increases.
Last, we found that electron affinity (exothermic reaction) is the energy paid (negative of the energy that occurs) when an electron is accepted by an atom in the gaseous state to form an anion (a negatively charged ion). A large electron affinity means that the atom wants the electron, so when looking at the periodic table, we find that there is an increase in the tendency to accept electrons from the left (halogens) to the right (metals), with the exception of the noble gases, which are stable (no affinity to lose or gain electrons).
Supplemental Material
If you would like to learn more on this topic, these resources may help:
The Periodic Table: Crash Course Chemistry #4
Completion Checklist
Make sure you complete the following learning activities and assessments before moving on to the next topic:
- Lecture Videos
- Example Videos (where posted)
- Practice Problems and Discussion
- Quiz