The Periodic Table Chapter 6

Periodic Table Arrangement of elements in which the elements are separated into groups based on a set of repeating properties Allows you to easily compare the properties of one element (or groups of elements) to another element (or groups of elements) Period: horizontal row Group (family): vertical column Number and letter (Group 1A)

Organizing the elements Mendeleev published his version of periodic table in 1869 He arranged the elements in order of increasing atomic mass. Figure 6.3 page 156 He left spaces predicting that elements would be discovered to fill in the holes and he predicted properties based on location on periodic table. Now elements are arranged by increasing atomic number

Periodic law The pattern of properties within a period repeats as you move from one period to the next

Metals Good conductors of heat/electricity High luster/sheen Reflects light Solid at room temp (except Hg) Ductile (pulled into wires) Malleable (hammered into thin sheets) Alkali metals: Group 1A (soft, explode in water) Alkaline earth metals: Group 2A (very reactive) Transition metals: B groups

Nonmetals Upper right hand corner Most are gases at room temperature Poor conductors heat/electricity (C is exception) Brittle (shattered when hit with hammer) Halogens (salt-formers): Group 7A (all 3 states of matter) Noble gases: Group 8A (inert gases)

Metalloids Border the stair-step Properties similar to metals and nonmetals Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Astatine

Coloring the Periodic Table Alkali metals: red Alkaline earth metals: dark orange Transition metals: light orange Inner transition metals: yellow Other metals: light pink Halogens: dark purple Noble gases: blue Other nonmetals: light purple Outline stair step: black Metalloids: green Key of colors

Metallic characteristics Metals want to lose electrons, not gain As you move from left to right, atoms tend to gain electrons, therefore, decreasing the metallic characteristics. Metals are to the left of the staircase and nonmetals are to the right of the staircase.

Atomic Size ½ the distance between the nuclei of two atoms of the same element when the atoms are joined Nuclear charge: strength of the nucleus’ pull Based on the number of protons in the nucleus Carbon as a nuclear charge of 6+ Shielding effect: electrons between outer electrons and nucleus shield the outer electrons from the full nuclear charge of the nucleus

Group trends in atomic size As atomic # increases within group, the charge on the nucleus increases and number of occupied energy levels increases. Increase in positive charge draws electrons closer to nucleus. Shielding effect is greater than the effect of the increase in nuclear charge Atomic size increases

Periodic trends in atomic size Across a period, electrons are added to the same principal energy level Shielding effect is constant for all the elements in the period Increasing nuclear charge pulls electrons in highest energy levels closer to nucleus Atomic size decreases

Ions Positive and negative ions form when electrons are transferred between atoms Cation: ion with positive charge Metals Loses electrons Na1+ Atom is larger than cation Anion: ion with negative charge Nonmetals Gains electrons Cl1- Atom is smaller than anion

Ionic Radius Radius of the ion. Cations are smaller than the atom they form from. Anions are larger than the atom they form from.

Ionization Energy Energy required to remove an electron from an atom Gaseous state First ionization energy is energy required to remove 1st electron Second ionization energy is energy required to remove the second electron

Electronegativity Ability of an atom to attract electrons when the atom is in a compound Noble gases do not have electronegativity because they do not form compounds Table 6.2 page 177

Electron Affinity Change in energy of a neutral atom in a gaseous state when an electron is added to form a negative ion Likelihood of an atom gaining an electron