1. Periodic Table Review

2. The placement or location of elements on the periodic table gives an indication of physical and chemical properties of that element. The elem ents on the periodic table are arranged in order of increasing atomic number. A few elements have smaller atomic masses than the elements that precede them (K, Ni, I). Be alert for a question that compares atomic masses of consecutive elements.

3. Elements can be classified as metals, nonmetals, metalloids, and noble gases. Know all the ways to distinguish between metals and non-metals. Metalloids rarely appear on the regents. You need to be able to identify them (on the step ladder) and know that they have some properties of metals and non-metals.

4. Elements can be differentiated by their physical properties. Physical properties of substances, such as density, conductivity, malleability, solubility, and hardness, differ between elements. Metals: very dense, conductors, malleable, usually hard Non-metals: not very dense (often gases), poor conductors, soft or brittle solids

5. Elements can be differentiated by chemical properties. Chemical properties describe how an element behaves during a chemical reaction. Metals lose electrons (+ oxidation #s) Non-metals gain electrons (- oxidation #s)

6. Some elements exist as allotropes. Allotropes are two or more forms of the same element that differ in their molecular or crystalline structure, and hence in their properties. The two examples they use are diamond/graphite for carbon, and ozone(O 3 ) and O 2.

7. For groups 1,2 and 13-18, elements within the same group have the same number of valence electrons (helium is an exception) and therefore similar reactivity. They often ask questions about elements with similar properties. The answer usually has two elements in the same group. Be able to draw Lewis diagrams for any element (# dots = # valence electrons)

9. The succession of elements within the same group demonstrates characteristic trends,e.g., differences in atomic radius, ionic radius, electronegativity, first ionization energy. These trends can be explained in terms of atomic structure(size).

10. Atomic radius increases down the group (larger size) Ionic radius also increases down the group (larger size) Electronegativity decreases down the group (elements become more metallic) Know the definition for electronegativity!! Ionization energy decreases down the group (valence electrons are shielded from the nucleus by lower energy levels)

11. The succession of elements across the period demonstrates characteristic trends, e.g., differences in atomic radius, ionic radius, electronegativity, first ionization energy. These trends can be explained in terms of atomic structure (nuclear charge).

12. Atomic Radius Atomic radius decreases across the period (more effective nuclear charge). Electrons from lower periods can “block” the nucleus from the valence electrons. The number of blockers does not change across the period, but the nuclear charge does!

14. Ionic Radius Ionic radius depends on # valence electrons: Non-metal ions are negatively charged, and are larger than their atomic radii. Metal ions are positively charged, and are smaller than their atomic radii. This is a common question on the exam!!!

16. Electronegativity Electronegativity increases across periods. Non-metals attract electrons more than metals. Electronegativity decreases down groups. Larger atoms have more electron shells, which shield the nucleus from available electrons.

17. Electronegativity of the Elements

18. Electronegativity Within Groups

19. Ionization Energy 1 st ionization energy generally increases across the period. More effective nuclear charge makes for smaller radii; it takes more energy to remove an electron. Ionization energy decreases down groups; valence electrons are further from the nucleus and shielded from it by electrons in filled energy levels