Mastering Periodic Table: Classification & Trends

1. Chapter 14 Chemical Periodicity

2. Section 14.1 Classification of the Elements • Objectives • Explain why you can infer the properties of an element based on those of other elements in the periodic table. • Use electron configuration to classify elements as noble gases, representative elements, transition metals, or inner transition metals.

3. Mendeleev’s Table • Mendeleev was the 1st to organize the elements • His method was according to atomic mass • The discovery of the atom’s insides gave way to a new organization of the elements

4. Old School Chemistry

5. The Periodic Table • The periodic table is one of the most important tools in Chemistry • Can predict properties • Tells you shapes and electron configurations • Shows sizes and other properties • It’s all about the electrons • The electron plays the most significant role in determining the physical and chemical properties of an element.

6. Element Classification • Elements can be divided into 4 groups based on their electron configurations • Noble Gases (Group 18) • Representative elements (s & p Blocks) • Transition metals (d Block) • Inner transition metals (f Block)

7. Noble Gases • Noble gases are elements in which the outermost s and p sublevels are filled. • Group 18 (or group 0) • Once called the inert gases. Inert means they don’t react, but under special circumstances, some noble gases do.

8. Representative Elements • In these elements, the outermost s and p sublevel is only partially filled. • These are usually called Group A elements. • alkali metals, alkaline earth metals, and halogens • The number of electrons in outer energy level is according to group number (group 3A has 3, etc.)

9. Transition Metals • These are metallic elements in which the outermost s sublevel and nearby d sublevel contain electrons. • These are all metals and are called Group B elements

10. Inner Transition Metals • These are metallic elements in which the outermost s sublevel and nearby f sublevel generally contain electrons. • Contain lanthanide and actinide series

12. Ex #1: Write the electron configurations of these elements: • Carbon 1s22s22p2 • Vanadium [Ar] 4s23d3 • Strontium [Kr] 5s2

13. Ex #2: What elements end with the following outer configurations? Classify these elements as noble gases, representative elements, transition metals, or inner transition metals. • s2 • Group 2 (2A – Alkaline Earth Metals) • s2p5 • Group 17 (7A – Halogens) • S2d2 • Group 4 (4B)

14. Table Review • Periods are principal energy levels (distance from nucleus) • Moving to the right of the periodic table, the next group will have 1 more electron in the sublevel than the previous group. • The d block in 3rd energy level is filled before p block in 4th energy level • Driving force for reactions is that they like their principal energy level to be full

15. Did We Meet Our Objectives? • Objectives • Explain why you can infer the properties of an element based on those of other elements in the periodic table. • Use electron configuration to classify elements as noble gases, representative elements, transition metals, or inner transition metals.

16. Section 14.2 Periodic Trends • Objectives: • Interpret group trends in atomic radii, ionic radii, ionization energies, and electronegativities. • Interpret periodic trends in atomic radii, ionic radii, ionization energies, and electronegativities.

17. Trends in Atomic Size • Sizes of atoms are not clearly known • Quantum mechanical model: edge of atom is fuzzy • The radius of an atom cannot be measured directly from one atom.

18. Periodic Trends: Atomic Size • Several ways to estimate sizes of elements • If a crystal, X-ray diffraction. • If a diatomic molecule, the distance between the nuclei can be estimated.

19. Periodic Trends: Atomic Size • Atomic radii – is one-half the distance between the nuclei of two like atoms in a diatomic molecule. • Atomic radii can be predicted by thinking about the number of electrons and forces between nucleus and electrons (opposites attract)

20. Decreases Increases

21. Periodic Trends: Atomic Size • Group Trends: Size increases moving down groups • The higher energy levels’ increasing distance from the nucleus decreases the attraction • Periodic Trends: Size decreases moving from left to right • Higher charge in nucleus will increase the attraction on electrons

22. Periodic Trends: Ionization Energy • Ionization energy: the energy required to remove outermost electron from a gaseous atom. • This can also be predicted

24. Ionization Energy • Group Trends: Energy decreases moving down groups. • Electron becomes farther from nucleus • Periodic Trends: Energy increases moving from left to right • Nuclear charge increases, increasing attraction

25. Periodic Trends: Ionization Energy • The energy required to remove the first outermost electron is called the first ionization energy. • Energy for removing 2nd electron is 2nd ionization energy and so on…. • Ionization energy increases as more electrons are removed

26. Periodic Trends: Ionic Size • Metals form cations easily • The cation is always smaller than neutral atom, loses electron(s) • Less repulsion between electrons • Nonmetals form anions easily • Anion is always larger than neutral atom, gains electron(s) • More repulsion between electrons

27. Periodic Trends: Ionic Size

28. Periodic Trends: Electronegativity • Electronegativity of an element is the tendency for the atom to attract electrons when combined with another element • Noble gases are not included • Transition metals don’t have a trend • These values have been calculated for the elements and are expressed in arbitrary units on the Pauling electronegativity scale. • These values will determine the type of bonding

31. Periodic Trends: Electronegativity

32. Pauling Electronegativity Values

33. Electronegativity Summary Electronegativity

34. Did We Meet Our Objectives? • Objectives: • Interpret group trends in atomic radii, ionic radii, ionization energies, and electronegativities. • Interpret periodic trends in atomic radii, ionic radii, ionization energies, and electronegativities.