Understanding Atomic and Ionic Radii
Chemistry classes spend a lot of time on periodic trends, and Chapter 10 Transparency 21 is one of those reference sheets teachers hand out when they want you to memorize how atoms compare in size. The topic covers atomic radius versus ionic radius and why they change across the periodic table. It is not complicated once you get past the initial confusion about why cations shrink and anions grow. The short answer key typically lists trends: atomic radius decreases moving left to right across a period and increases moving down a group. Ionic radius follows similar logic but with an important twist depending on whether the ion is positive or negative. The transparency usually includes a chart comparing sodium to sodium ion, chlorine to chloride, and a few other common examples so students can see the numbers directly.
Chapter 10 Transparency 21 Atomic And Ionic Radii Answers
Here is the practical breakdown most answer keys will have for that transparency. Atomic radius gets smaller across a period because protons are added to the nucleus without adding a new electron shell. The increased nuclear pull draws the electron cloud closer. Moving down a group adds electron shells, which makes the atom larger regardless of the proton count. Typical values for period 3 elements run from about 186 picometers for sodium down to roughly 100 picometers for chlorine. When an atom loses electrons to become a cation, the remaining electrons are held tighter by the same number of protons. The ion shrinks. Sodium loses one electron to become Na, and its radius drops from about 186 pm to roughly 102 pm. When an atom gains electrons to become an anion, electron-electron repulsion pushes the cloud outward. Chlorine gains an electron to become Cl, and its radius expands from about 99 pm to roughly 181 pm. These are the numbers the transparency is showing.
I remember grading these transparencies back when I was tutoring introductory chemistry. The problem students consistently struggled with was confusing which direction the trend went for isoelectronic species. Three ions like O², F, and Na all have ten electrons, but their sizes differ because the nuclear charge varies. More protons pulling on the same number of electrons means a smaller radius. Students kept picking the heaviest element as the smallest without thinking through the charge balance. I started having them write out the proton count above each symbol before they answered anything. That simple step cut the error rate dramatically over a semester. One thing most textbooks gloss over is that ionic radius values are not fixed constants. They depend on coordination number and the method used to measure them. Pauling radii, Schomaker-Stevenson radii, and Goldschmidt radii can give slightly different numbers for the same ion. If you are doing calculations that require precision, you should note which scale your source uses. For a high school transparency exercise this does not matter much. The relative ordering stays the same across all scales. Another counter-intuitive point is that transition metals do not follow the same clean trend. Their ionic radii change very little across the series because the d-electrons shield each other poorly but also do not add new shells. You might see a lanthanide contraction effect where period 6 transition metals end up nearly the same size as period 5 ones. That is why hafnium and zirconium have almost identical radii despite being two periods apart. This usually does not appear on a basic transparency sheet, but it shows up in more advanced courses if you connect the dots later.
If you need to download or print the actual transparency for classroom use, check your publisher's resources page or ask your teacher for the digital copy. Most Glencoe or McGraw-Hill chemistry courses include these as downloadable PDFs in the teacher resource folder. The answer key itself is typically separate and not always posted publicly. If you are stuck on a specific question from that transparency, the trend logic above will cover most of it. The main pitfall to avoid is assuming that a larger atom always forms a larger ion. Magnesium is smaller than sodium, but Mg² is also significantly smaller than Na. The charge matters as much as the position on the table. Similarly, aluminum forms Al³ which is even smaller still. Keep the proton-to-electron ratio in mind and the answers become straightforward.
