The Periodic Table Unit 2

Periodic Law Periods = horizontal rows Seven periods in table. Each period corresponds to a principal energy level. Groups/Families = vertical columns Within a group, the properties of the elements change as you move across a period but repeat as you progress down the column. Periodic law: when elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

Metals, Nonmetals & Metalloids Elements can be grouped into three broad classes based on their general properties. Three classes of elements are metals, nonmetals, and metalloids. Across a period, the properties of elements become less metallic and more nonmetallic.

Metals About 80% of elements are metals. Metals – good conductors of heat & electric current. All metals are solids at room temperature except mercury. Common Metallic Properties Exhibit high luster or sheen on a clean or cut surface. High melting points & high densities Large atomic radii Low ionization energies & low electronegativities Malleable & ductile

Nonmetals Common nonmetal properties Nonmetals are poor conductors of heat and electric current. Carbon is an exception. More variation in the physical properties of nonmetals than with metals. Most are gases at room temperature, a few are solids, and one, bromine, is a liquid. Common nonmetal properties Solid nonmetals are brittle. High ionization energy & high electronegativity. Little or no metallic luster Gain electrons easily

Metalloids Common metalloid properties Most of the elements that border the stair step line separating metals from nonmetals. A metalloid generally has properties that are similar to those of metals and nonmetals. Depending on the conditions, a metalloid can behave as a metal or a nonmetal. Common metalloid properties Electronegativities & ionization energies between metal and nonmetal. Reactivity depends on other elements in reaction Often make good semiconductors.

Classifying the Elements The periodic table displays the symbols and names of the elements, along with information about the structure of their atoms. Group 1 or 1A elements are called alkali metals. Group 2 or 2A elements are called alkaline earth metals. Group 17 or 7A are called halogens. Group 18 or 8A are called noble gases. Sometimes called inert gases because they rarely react. Completely filled s and p levels.

Representative Elements Portion of periodic table containing Groups 1 or 1A through 2 or 2A and also 13-17 or 3A - 7A. Elements in these groups are referred to as representative elements because they display a wide range of physical and chemical properties. Some are metals, some are nonmetals, and some are metalloids. Most are solid, a few are gases and one is a liquid.

Electron Configurations in Groups Electrons play key role in determining the properties of the elements. Remember that protons give an element its identity, electrons determine an elements reactivity. Elements can be sorted into noble gases, representative elements, transition metals, or inner transition metals, based on their electron configurations.

Transition Elements Group B Elements – separate both sets of Group A elements (1-2 and 13-18). Group numbers 3-12 when table is numbered 1-18. Two types of transition elements - transition metals and inner transition metals. Classified based on electron configurations. In atoms of transition metals, the highest occupied s sublevel and a nearby d sublevel contain electrons. These elements are characterized by the presence of electrons in d orbitals.

Transition & Inner Transition Metals Low ionization energies Positive oxidation states Very hard High melting points High boiling points High electrical conductivity Malleable When progressing from left to right on periodic table, the five d orbitals become filled. The highest occupied s sublevel and a nearby f sublevel generally contain electrons. Inner transition metals are classified by f orbitals containing electrons.

Electron Orbitals

Sublevel Blocks Electron configuration, as shown in the previous slide, is another factor in the placement of elements on the periodic table. s block = Groups 1A, 2A, and helium. p block = groups 3A, 4A, 5A, 6A, 7A & 8A except helium. d block = transition metals f block = inner transition metals

Periodic Trends - Atomic Size Periodic trends are trends that exist in the periodic table can be explained by variations in atomic structure. Atomic radius – one half the distance between the nuclei of two joined atoms of the same element. In general, atomic size increases from top to bottom within a group and decreases from left to right across a period. See next slide for example.

Atomic Radii Trend

Atomic Size Group Trend Atomic # increases  charge on nucleus increases  # of occupied energy levels increase Affect atomic size in opposite ways Increase in nuclear charge draws electrons closer to nucleus. Increase in energy levels shields electrons in outer energy levels from the nuclear attraction. Shielding effect > effect of nuclear charge = increase in atomic size.

Atomic Size Period Trend From left to right across a period, each element has 1 more proton and electron than the element before it. Across a period, elements are added to the same principal energy level. Shielding effect is constant for all elements in period. Increasing nuclear charge pulls electrons closer causing atomic size to decrease.

IONS Ion – an atom or group of atoms that has a positive or negative charge. Ions are formed when electrons are transferred between atoms. Metals tend to lose electrons forming positive ions. Nonmetals tend to gain electrons forming negative ions. Cation – an ion with a positive charge. Charge is written as a number followed by a plus sign. i.e.: Na1+ or simply Na+ where the one is assumed. Anion – an ion with a negative charge. Charge is written as a number followed by a negative sign. i.e.: Cl1- or simply Cl- where the one is assumed.

Trends in Ionic Size Cations are always smaller than the atoms from which they form. Anions are always larger than the atoms from which they form.

Trends in Ionization Energy Ionization energy – the energy required to remove an electron from an atom. 1st ionization energy – energy required to move 1 electron from an atom. 1st ionization energy tends to decrease from top to bottom within a group and increase from left to right across a period. As size of atom increases, nuclear charge has less effect on outermost electrons, resulting in less energy required to remove an electron. Less energy required to remove an electron the lower the 1st ionization energy.

Periodic Trends in Ionization Energy In general, 1st ionization energy of representative elements increases from left to right across a period. Reason: Increasing nuclear charge + constant shielding effect = greater attraction of nucleus for electrons. Result : It takes more energy to remove an electron. The smaller the ion, the stronger the ionization energy.

Ionization Energy Trend

Trends in Electronegativity Electronegativity is a property that can be used to predict what type of bond will form during a reaction. Electronegativity – the ability of an atom of an element to attract electrons when the atom is in a compound. In general, electronegativity decreases from top to bottom within a group. For representative elements, the values tend to increase from left to right across a period.