1. Chapter 6 The Periodic Table and Periodic Law

2. Development of the Modern Periodic Table Modern Periodic Table Periodic law – states that there is a periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number Arranged in order of increasing atomic number into a series of columns, called groups, and rows, called periods Groups are labeled 1-8, followed by A and B “A” groups are the representative elements b/c they possess a wide range of chemical and physical properties “B” groups are the transition elements

3. Classifying elements – 3 main classifications for elements Metals Located on left side of periodic table (H is exception) Shiny, solid at room temp., good conductors, malleable, ductile Group 1A (except H) – alkali metals oVery reactive Group 2A – alkaline earth metals oReactive Transition elements Located in the middle section of the periodic table oTransition metals oInner transition metals – very bottom of the table to save space Lanthanide series Actinide series

4. Nonmetals Located in upper right side of table Generally gases or brittle, dull solids, poor conductors oOnly bromine is liquid at room temp. Group 7A – halogens oHighly reactive Group 8A – noble gases oUnreactive Metalloids Located on border of stair-step line Physical and chemical properties of both metals and nonmetals

5. Classification of the Elements Organizing the elements by electron configuration Valence electrons – one of the most important relationships in chemistry Atoms in the same group have similar chemical properties b/c they have the same number of valence electrons Valence electrons and period – the energy level of an element’s valence electrons indicates its period Example: Li’s valence electron is in 2 nd energy level and Li is in period 2; Ga’s electron configuration is [Ar]4s 2 3d 10 4p 1, its valence electrons are in the fourth energy level, and it’s found in the 4 th period Valence electrons and group number – representative element’s group number and valence electrons are related Noble gases have 8 valence electrons (except He)

6. The s-, p-, d-, and f-block elements – b/c there are 4 different energy sublevels (s, p, d, f), the periodic table is divided into 4 distinct blocks (Fig 6-10) S-block – groups 1A, 2A, and H and He P-block – groups 3A – 8A Group 8A (noble gases) are unique b/c of their stability; they undergo virtually no chemical reactions D-block – transition metals F-block – inner transition metals Period patterns Period 1 contains only s-block Periods 2 and 3 contain s- and p-block Periods 4 and 5 contain s-, p-, and d-block Periods 6 and 7 contain s-, p-, d-, and f-block

7. Periodic Trends Atomic radius Trends within periods – general decrease in atomic radii left-to- right No additional electrons come between the valence electrons and nucleus Trends within groups – general increase in atomic radii moving down Outermost orbital increases in size along w/ increasing principal energy level, making atom larger Fig 6-12

8. Ionic radius Atoms can gain or lose electrons to form ions Ion – atom or bonded group of atoms that has a positive or negative charge When atoms gain electrons and form negatively charged ions, they always become larger When atoms lose electrons and form positively charged ions, they always become smaller Ionization energy – energy required to remove an electron from a gaseous atom First ionization energy - energy required to remove the first electron from an atom Indication of how strongly an atom’s nucleus holds onto its valence electrons High ionization energy – atom has strong hold on electrons and are less likely to form positive ions

9. Trends within periods and groups Fig 6-17 Octet rule – atoms tend to gain, lose, or share electrons in order to acquire a full set of 8 valence electrons Elements on right side of periodic table tend to gain electrons Form negative ions Elements on left side of table tend to lose electrons Form positive ions Electronegativity – indicates the relative ability of its atoms to attract electrons in a chemical bond Fig 6-18