Periodic Properties of Elements

Atomic Radius Atomic radius – the distance from the nucleus to the outermost electrons, measured in picometers ( m)

Ionization Energy Ionization Energy - the energy needed to remove the outermost electron Elements have multiple ionization energies: 1 st, 2 nd, 3 rd, etc.

Lab 5: Periodic Properties of Groups of Elements Objective To examine the periodic properties of elements in groups (columns) of the periodic table. Materials Chemistry Reference Tables, graph paper, rulers, colored pencils.

Table S

Procedure Using the blank partial periodic table provided, write the symbol of the element, the sublevel ground state electron configuration, the atomic radius of the element, and the first ionization energy of the element for the following: Group 1: the alkali metals (except hydrogen) Group 2: the alkali earths Group 17: the halogens Group 18: the noble gases (except helium)

3 Li 1s 2 2s 1 Radius ______pm IE ________kJ

The atomic radius and first ionization energy data can be found on Table S. Complete the table and follow the directions below for making two graphs: one for atomic radius of the four groups, the other for 1 st ionization energy.

Making Graphs You will be making two multiple line graphs that each contain four sets of data (four lines). Each group should be treated as a separate line; the line must be drawn in a different color. Plot the atomic number, from 1-90, on the x axis. On the y axis, plot the atomic radius, in picometers. Make a suitable scale for each axis so that each data point can fit on the graph. Be sure to indicate which group is which on your graph. You can use a legend for this, or you can label the individual lines. The graph must have a descriptive title. For your second graph, plot 1 st ionization energy (in kJ) on the y axis. The x axis should be similar to the first graph. This graph should be completed as homework for the next day if not finished in class.

Questions (answer in complete sentences, on a separate sheet of paper) 1. What happens to the atomic radius as you go down a group? Give a reasonable explanation for this trend. 2. What happens to the ionization energy as you go down each group? Account for this phenomenon. 3. Using sublevel electron configuration, explain why elements within groups have similar physical and chemical properties. 4. Define the following terms: a) dependent variable b) independent variable 5. For each graph, identify which axis is the dependent and independent variable, respectively.

Lab 6: Periodic Properties Within Energy Levels of the Periodic Table Objective: to examine the periodic properties of elements in periods 2 and 3 of the periodic table. Same materials as lab 5.

Methods Using the blank partial periodic table provided, write the symbol of the element ground state electron configuration the atomic radius of the element first ionization energy of the element for periods 2, 3, and the first two elements of period 4. Some of the data will be repeated from your previous assignment; copy that data onto this sheet.

Making Graphs On the x axis, plot the atomic number, from On the y axis, plot the atomic radius, using a suitable scale. Units should be included. Plot your points, putting an appropriate title on the graph. On your graph, you can identify the two periods and/or identify each element at every point if you wish. Repeat the same process for ionization energy. The assignment must be completed by the following class as homework.

Questions (answer in complete sentences, on a separate sheet of paper) 1. What happens to the atomic radius as the atomic number increases across a period? Give a reasonable explanation for this trend. 2. In general, what happens to the ionization energy as the atomic number increases across a period? Explain your answer. 3. Compare ionization energies for the metallic elements to the non-metals. Why are they so different? 4. Identify the dependent and independent variables for each graph.

Extra credit Filled sublevels and half filled sublevels have extra stability. Elements that do not have these configurations have electrons in somewhat higher energy states and are easier to ionize. Using this information and your knowledge of electron configurations, explain why there are two little “dips” in the 1 st ionization energy as you move across each period.

Atomic Radius

Atomic Radius increases down groups Atomic Radius decreases across periods

Atomic Radius increases down groups Why? More energy levels with larger elements Higher energy levels, further from the nucleus, are filled with electrons

Atomic Radius decreases across periods Why? The nuclear charge is increasing across the period The number of “blocking” electrons does not change – the blockers are in the lower energy levels Electrons in the highest energy level don’t shield each other from the nucleus The charge on consecutive elements becomes more effective, drawing the electrons closer to the nucleus.

Ionic Radius

Metals tend to have large atomic radii, and small ionic radii When metals lose electrons, they now have more protons than electrons, and the remaining e - are pulled closer to the nucleus Non-metals tend to have small atomic radii, and large ionic radii When non-metals gain electrons, their increased repulsion enlarges the atom, and the nuclear charge is less able to keep the electrons close

Atomic and Ionic Radii Questions Which is larger, Na or Na + ? Which is larger, Cl or Cl - ? Which is larger, Li or F? Which is larger, K or Rb? Which is larger, F or Br? Which is larger, Na + or K + ? Which is larger, I - or Br - ?

Ionization Energy Decreases down groups Generally increases across the table

Periodic Trends for Ionization Energy For elements of the same group, ionization energy decreases due to shielding by more and more energy levels filled with electrons. The valence electrons, being at higher energy than their cousins in lower periods, require less energy to leave the atom.

Ionization Energy Across Periods Ionization generally increases across periods, because the effective nuclear charge is increasing with each added proton. There a few exceptions due to interactions between electrons in groups 13 and 16.