TABLE PERIODIC

Development of the Periodic Table Newlands – law of octaves Mendeleev – the periodic table

Periodic Law Elements in order of increasing atomic number Fe Co Ni Cu Atomic no. above Atomic mass below

Periodic Law Repeating pattern of characteristics (periodicity)

Lewis dot diagrams Valence shell electrons Atomic Structure and Periodicity Lewis dot diagrams Valence shell electrons 1 2 13 14 15 16 17 18

Large Groups Metals, Nonmetals, Metalloids Transition Metals Rare Earth Elements

READING THE TABLE Groups (Families) Cu, Ag, Au List 5 characteristics that Cu, Ag, and Au have in common

READING THE TABLE Periods – 1, 2, 3… What characteristics do the elements in period 1 all have in common?

Metals Left side of periodic table, below staircase Luster - shine Good conductors of heat and electricity Solids (room temp), exception-mercury Malleable-can be hammered into thin sheets Ductile-can be drawn into fine wires Lose electrons in bonds Make positive ions

Nonmetals Far right side of periodic table above staircase Poor conductors of heat and electricity Neither malleable nor ductile Gases or solid (room temp), 1 liquid (bromine) Gain electrons in bonds Make negative ions

Semimetals Metalloids ex. Si, As, Ge, Sb, Te Both properties of metals and nonmetals Form + ions Semiconductors

Alkali Metal Li, Na, K … Shiny solids, malleable, ductile, good conductors Low densities, low melting point Very reactive (Na most active) Soft enough to be cut by a knife Intense reactions with water and air One valence electron Make 1+ ions Na, K most abundant of group, always found in nature in compounds (due to reactivity) Na: baking soda, table salt, bleach

Alkaline Earth Metals Be, Mg, Ca, Sr, Ba, Ra Densities and melting points > group 1A 2 Valence electrons 2+ charge on ions Not found in nature as elements Mg and Ca most abundant in group Mg wheels, tools CaCO3 shells, chalk

Halogens F, Cl, Br, I “Salt Former” All exist in elemental forms as Diatomic Molecules F2 Cl2 Br2 I2 High reactivity (React with most metals and nonmetals) 7 valence electrons Make 1- ions F corrosive gas, most reactive element CFC refrigerators, air conditioners Cl pool cleaner, PVC plastic, bleach Br pesticides, photo film I iodized salt, alcohol solution for cuts

Noble Gases He, Ne, Ar, Kr, Xe, Rn Least reactive of all elements 8 valence electrons Do not make ions Ar light bulbs Ne signs He balloons Rn is a radioactive gas

Transition Metals Iron and Titanium, abundant; Platinum, rare High densities, high melting points Make alloys Relatively non-reactive Make colored compounds Luster Good conductors of electricity and heat Make variable ions

Inner Transition Metals Lanthanides (begins with La) -soft, silvery metals, tarnish in air -making steel alloys Actinides (begins with Ac) -all isotopes are radioactive -nuclear fuels

Hydrogen Nonmetal, colorless, odorless 1 valence electron 1+ charge Elemental H2 Most is combined with oxygen as water Most abundant element in universe Major use--making ammonia for fertilizer

PERIODIC TRENDS Atomic Radius Ionization Energy Electronegativity Ion Size Electron Affinity

Atomic Radius (pm) Plotted against Atomic Number

Atomic Radius Distance from the center of the nucleus to the outside shell of electrons

Atomic Radius Why does atomic radius increase as you go down a column or family? Why does atomic radius decrease as you go across a row or period?

Practice Problems K Al P Cl As S F He Ra Mg Be Ca F B O N Li Al Sb At

Atomic radius increases as you go down a group because you are adding more electron shells/ adding more energy levels. Atomic radius decreases across a period because you are adding more protons, increasing the effective nuclear charge and the electrons are being added to the same energy level. The nucleus pulls that electron shell in tighter

Ionization Energy

Ionization Energy Energy needed to remove an e- in gas state Tells you how strongly an atom holds onto its valence electron Unit-joules (J or kJ)

Ionization Energy

Ionization energy increases across a period because the valence electrons are closer and closer to 8. High IE means they don’t want to give up electrons. Low IE means they don’t mind giving up electrons. Ionization energy decreases down a group because of “shielding” – they have more electron shells, further from the nucleus (+), so they are easier to pull away

Successive Ionization Energies Mg  Mg+ + e- 735 kJ/mol Mg+  Mg2+ + e- 1445 kJ/mol Mg2+ Mg3+ + e- 7730 kJ/mol

Successive Ionization Energies Al  Al+ + e- 580 kJ/mol Al+  Al2+ + e- 1,815 kJ/mol Al2+ Al3+ + e- 2,740 kJ/mol Al3+ Al4+ + e- 11,600 kJ/mol

Practice Problems K Al P Cl As S F He Ra Mg Be Ca F B O N Li Al Sb At

Electronegativity

Electronegativity The attraction that an atom has for electrons in a chemical bond F is 4 (highest) Fr is lowest

Practice Problems K Al P Cl As S F He Ra Mg Be Ca F B O N Li Al Sb At

Ionic Size Cations (Positive Ions) - smaller than atom A more positive nucleus pulls electron shells in tighter Anions (Negative Ions) - larger than atom More electrons= more resistance to positive nucleus

Ionic Radii

Atomic Radius Ionic Radii Electronegativity Whichever element is closest to have 8 electrons in its valence shell has the highest electronegativity. The closer an element has to 8 electrons in its valence shell, the it wants to attract electrons. Noble gases have a full valence shell and DO NOT want to attract electrons. They have the lowest electronegativity For elements in the same group, the element furthest down has the most electron shells, so the valence shell is less affected by the attraction of the nucleus (lowest electronegativity) Ionization Energy Whichever element is closest to have 8 electrons in its valence shell has the highest ionization energy. The closer an element has to 8 electrons in its valence shell, the more difficult it is to pull an electron away. For elements in the same group, the element furthest down has the most electron shells, so the valence shell is further from the nucleus, making it easier to pull an electron away (lowest ionization energy) Atomic Radius Whichever element is furthest down has the most energy levels/electron shells (has the largest radius) For elements in the same period, the element furthest to the right has the highest effective nuclear charge, which pulls the electrons in tighter (has the smallest radius) Ionic Radii Whichever ion is the most negative has the largest atomic radius because it is less affected by the positive nucleus. Whichever ion is the most positive has the smallest radius because it has a higher effective nuclear charge.