Ch. 6 The Periodic Table

6.1 Organizing the Elements Early chemists used properties of elements to sort them into groups J.W. Dobereiner classified elements into triads with similar properties Noted that 1 of the three had properties (i.e. mass) in between the other two Ex: average of the masses of iodine and chlorine is very near the mass of bromine

Mendeleev’s Periodic Table Developed for students Arranged elements in order of increasing atomic mass* Periodic repetition of properties Left blank spaces where no known element fit the mass and properties Elements later discovered that were predicted by this table

Today’s Periodic Table Arranged in order of increasing atomic number Periodic Law: when elements are arranged by increasing atomic number, there is a periodic repetition of physical and chemical properties

Metals, Nonmetals, and Metalloids Three broad classes of elements

Metals Good conductors of heat and electrical current Luster/sheen: ability to reflect light Solid at room temperature (except Mercury) Ductile: can be drawn into wires Malleable: can be hammered into sheets without breaking

Nonmetals Tend to have properties opposite those of metals Poor conductors (except carbon) Brittle solids

Metalloids Properties similar to both metals and nonmetals under different circumstances

6.2 Classifying the Elements Group 1: alkali metals Group 2: alkaline earth metals Elements 57-70 and 89-102: inner transition metals Groups 3-12: transition metals Metals in Groups 13-17: other metals Along zigzag line from Groups 13-17: metalloids Nonmetals in Groups 14-16: nonmetals Group 17*: halogens Group 18: noble gases

Grouping by Electron Configuration Noble Gases: highest energy level is completely full

Grouping by Electron Configuration Representative elements: s or p sublevel of the highest energy level is not filled Group number (3A, 4A, etc) is the same as the number of electrons in the highest energy level

Grouping by Electron Configuration Transition metals: electrons present in d orbitals Inner transition metals: electrons present in f orbitals

Write the electron configurations of A) carbon B) strontium C) vanadium List the symbols of all the elements with electron configurations ending in ns2 np1 ns2np5 nd2 n can be 2, 3, 4, 5, 6, 7 of the highest occupied energy level

Electron configuration of A) carbon: 1s22s22p2 B) strontium: 1s22s22p63s23p64s23d104p65s2 C) vanadium: 1s22s22p63s23p64s23d3 ns2 np1 B, Al, Ga, In, Tl, Uut ns2np5 F, Cl, Br, I, At, Uus nd2 Ti, Zr, Hf, Rf

6.3 Periodic Trends Trends in atomic size (atomic radius) Atoms increase in size from top to bottom within a group More protons, neutrons, and electrons take up more space, different energy levels, etc. Atoms decrease in size from left to right across a period As more protons and electrons are added (electrons in the same energy level) the increasing charge of the nucleus pulls electrons in closer Shielding effect: if far enough away, charge attraction is ineffective Atomic radius: ½ the distance between 2 nuclei of two atoms of the same element joined together

Ions Ions are formed when a neutral atom gains or loses electrons Positive ions (cations) are formed when electrons are lost Negative ions (anions) are formed when electrons are gained

Ionization Energy Ionization energy is the energy required to remove an electron from an atom The first ionization energy is the energy required to remove the first electron The second ionization energy is energy to remove a second electron Ionization energy decreases from top to bottom in a group Ionization energy increases from left to right across a period Subsequent ionization energies are greater than the first ionization energy

Ion Size Ionic size is different from the neutral atom size Cations are smaller than the neutral atom Anions are larger than the neutral atom Ionic size increases from top to bottom in a group Ionic size of anions and cations decreases from left to right in a period

Ion Size

Electronegativity Electronegativity: ability of an element to attract electrons to itself in a compound Can be used to predict the type of bond that will form between two elements Example: in a water molecule, oxygen is more electronegative The electrons in the compound are drawn closer to (or held more tightly by) the oxygen atom than the hydrogen atoms

Electronegativity Trends Electronegativity values decrease from top to bottom in a group Electronegativity values increase from left to right across a period Noble gases do not have electronegativity values

What Ion does it form? Li Mg Rb Ca K P Cl Be F O N I Al Br S

What Ion does it form? Li+1 Mg+2 Rb+1 Ca+2 K+1 P-3 Cl-1 Be+2 F-1 O-2 Al+3 Br-1 S-2

Ch. 6 Assessment Questions