The Periodic Table

Guiding Questions Why is the periodic table so important? Why is the periodic table shaped the way it's shaped? Why do elements combine? Why do elements react? What other patterns are there in the world and how do they help us? Periodic Table Study Questions   1.         Why did chemists make the periodic table? 2.         Why was the table difficult to make? 3.         Why were Dobereiner’s triads of limited use at a periodic table? 4.         What did Newland discover about the elements? 5.         What did Meyer contribute to the development of the periodic table? 6.         What did Mendeleev use as the organizing property for the periodic table? 7.         What problem developed from the use of this property? 8.         What is common to elements in a column of the table? 9.         How did properties change in a row of the table? 10.       What was the significance of gaps in Mendeleev’s periodic table? 11.       What did Moseley use to order the elements in the periodic table? 12.       How did Moseley change the periodic law? 13.       What determines the identity of an element? 14.       Why do elements in a column of the periodic table have similar properties? 15.       With respect to the Periodic Table, what is the meaning of periodic? 16.       What does a row of the Periodic table represent? 17.       What happens to valence electrons as you move left to right in a row? 18.       When determines stability in an atom? 19.       List, from least to most, the stable configurations in an atom. 20.       What determines the column of the periodic table an element is in? 21.       What sublevels are in the outer level of an atom? 22.       What is the maximum number of electrons in the outer level of an atom? 23.       What determines the row and column of the periodic table an element is in? 24.       What are common properties of metals? 25.       What are common properties of non-metals? 26.       What three things can happen to electrons when atoms form compounds? 27.       The configuration of He is 1s2, but it is placed in column 18. Explain this discrepancy. 28.       Hydrogen is obviously not an alkali metal. Why is it in column 1 of the table? 29.       What is necessary for a metalloid to act as a semiconductor?

Mendeleev “Father of Periodic Table” organized elements based on increasing atomic mass. Found similarities in chemical properties and published his first periodic table in 1869.

Moseley discovered while working with elements that they fit better into pattern when arranged by nuclear charge (number of protons-also known as atomic number).

Periodic Law states that physical and chemical properties of the elements are periodic functions of their atomic numbers. This means that when elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals

Organization of the Periodic Table Periods – rows across, horizontal Indicates number of energy levels (Principle Quantum Number) Groups/Families – down, columns 18 Indicate number of e- in outer most energy level (1-2, 13-18 main group elements) Share chemical properities

Three Categories of Elements Metals Non-Metals Metalloids

Three Categories of Elements Metals - Groups 1-12 (except H) and under stair-step groups 13-15 - Form ionic and metallic bonds Luster (shiny, reflects light) Malleable - can be flattened into sheets Ductile - can be drawn into thin wires Good Conductors - heat and electricity can flow throughbecause the outer electrons are not held tightly to the nucleus and move freely - Most are solid at room temperature except Mercury

The Periodic Table H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Rg The Metals are represented in the Periodic Table in blue.

Three Categories of Elements Non-Metals Dull Not Malleable/Ductile Poor Conductors - P block - Form ionic and covalent bonds - Most do not conduct heat or electricity -All, except H, are found on right of periodic table - At room temperature, most are solid or gaseous

The Periodic Table H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Rg The Non-Metals are represented in the Periodic Table in yellow.

Three Categories of Elements Non-Metals Some Examples of Non-Metals Oxygen (O) Helium (He) Sulfur (S) Chlorine (Cl) Neon (Ne) Nitrogen (N)

Three Categories of Elements Metalloids Have properties of both metals and non-metals. Generally not shiny Not malleable and ductile - Form ionic and covalent bonds They conduct heat and electricity better than nonmetals, but less than metals.

Three Categories of Elements Metalloids Here are the Metalloids Boron (B) Arsenic (As) Tellurium (Te) Silicon (Si) - Polonium (Po) Antimony (Sb) Germanium (Ge)

The Periodic Table H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Rg The Metalloids are represented in the Periodic Table in green.

The Periodic Table H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Rg What do you notice about the way the element groups are arranged in the Periodic Table?

Hydrogen In a class by itself Most common element in the universe (3/4) Behaves unlike any other element because it consists of one proton and one electron.

Main Group Elements S and p blocks Groups 1,2, and 13-18 Four groups have special names: Alkali metals (Group 1) Alkaline-earth metals (Group 2) Halogens (Group 17) Noble gases (Group 18)

Group 1 – Alkali Metals Characteristics Soft, silver metals Low melting points and densities Highly reactive (esp. w/ water) Do not occur in nature in elemental form Stored in kerosene or mineral oil Have one Valence electron

Group 2 – Alkaline-earth Metals Characteristics Gray, metallic solids Reactive; but, less than alkali Not found as free elements in nature Bright fireworks, aircraft Harder, denser, and stronger than alkali metals Higher melting points than alkali Have 2 valence electrons

Transition Elements Groups 3-12 (elements in transition) D block Vary in reactivity and can be found as free elements Metal characteristics Form colored compounds Often occur in nature as uncombined elements Hg – mercury – liquid metal

Valence Electrons for Transition Elements Group 3: 3 valence electrons Group 4: 2 to 4 valence electrons Group 5: 2 to 5 valence electrons Group 6: 2 to 6 valence electrons Group 7: 2 to 7 valence electrons Group 8: 2 or 3 valence electrons Group 9: 2 or 3 valence electrons Group 10: 2 or 3 valence electrons Group 11: 1 or 2 valence electrons Group 12: 2 valence electrons

Group 13 – Boron Group or Icosagens Have 3 valence electrons Boron is the only metalloid in the family The rest are poor metals.

Group 14 – Carbon group or Crystallogens Have 4 valence electrons Unique feature is that the elements can form different anions and cations Ex: C: 4- Si and Ge: 4+ Sn and Pb: 2+

Group 15 – Nitrogen Group or Pnictogens Have 5 valence electrons Able to form double and triple bonds

Group 16 – Oxygen group or Chalcogens Have 6 valence electrons From -2 ions Physical properities vary dramatically Ex: oxygen – colorless gas Sulfur – yellow solid Tellurium – silver metalloid Selenium - black

Group 17 - Halogens Most reactive non-metal Irritating odor Interact with alkali metals to form salts 7 electrons in outer energy level Easily pickup one electron Bromine – only liquid nonmetal Flourine and chlorine – gases at room temp Iodine and astatine – solids at room temp

Group 18 – Noble Gases No color or odor Exist as individual gas atoms (monatomic) Full outer energy level Relatively un-reactive (inert gases)

Inner Transitional Metals f-block, n-2 ALL are radioactive and unstable Includes lanthanides (atomic number 58-71) Shiny metals, similar in reactivity to alkaline earth metals and actinides (atomic number 90-103) First 4 are found on earth, the remaining are synthetic

Diatomic Molecules Elements That Exist as Diatomic Molecules in Their Elemental Forms Element Present Elemental State at 25 oC Molecule hydrogen colorless gas H2 nitrogen colorless gas N2 oxygen pale blue gas O2 fluorine pale yellow gas F2 chlorine pale green gas Cl2 bromine reddish-brown liquid Br2 iodine lustrous, dark purple solid I2

How do we use an elements location on the periodic table to determine its ionic charge? Atomic Radius (Angstrom) ½ the distance from the nuclei to another From the nucleus to edge of e- cloud Going down a group, Atomic radius increases because of the increasing number of energy levels Going across a period (left to right), atomic radius decreases because of the increase in positive charge in the nucleus

Atomic Radii

Coulombic attraction Attraction of + and – charges Two factors determine the strength: Amount of charge Distance between charges

Shielding Effect Kernel electrons “shield” valence electrons from attractive force of the nucleus Caused by kernel and valence electrons repelling each other The more electron shells there are, the greater the shielding effect. Explains why valence electrons are more easily removed

+ Shielding Effect - - - - nucleus Electron Valence + - nucleus - - Electrons - Some schools use the term “screening” as I use the term “shielding”. Electron Shield “kernel” electrons Kernel electrons block the attractive force of the nucleus from the valence electrons

Why is cesium bigger than sodium? Sodium has the electron configuration 1s22s22p63s1. There is one valence electron. The attraction between this lone valence electron and the nucleus with 11 protons is shielded by the other 10 core electrons. The electron configuration for cesium is 1s22s22p63s23p64s23d104p65s24d105p66s1. While there are more protons in a cesium atom, there are also more electrons shielding the outer electron from the nucleus. The outermost electron, 6s1, therefore, is held very loosely. Because of shielding, the nucleus has less control over this 6s1 electron than it does over a 3s1 electron.

Ionization Energy Ionization is a process that results in the formation of an ion An ion is an atom or group of atoms that have a “+” or “-” charge Change is created by gain or loss of e- Losing an e creates a “+” charge (cation)

(also known as 1st ionization energy) Losing the e- requires energy. Energy required to remove one e- from a neutral atom is called IONIZATION ENERGY (also known as 1st ionization energy)

Factors that affect ionization Nuclear charge – the larger the nuclear charge , the greater IE Shielding effect – the greater the shielding effect, the less IE Radius – greater the distance between the nucleus and the outer electrons of an atom, the less IE. Sublevel – an electron from a sublevel that is more than half-full requires additional energy to be removed

Ionization energy decreases as you go down a group because valence electrons are farther from the nucleus And increases as you go across a period because of the greater positive charge leads to greater attraction to electron

Ionization Energy

ELECTRON AFFINITY Electron Affinity Gaining an electron results in an ion with a “-” charge (anion) When an atom gains an e- it causes an energy change. The energy change when a neutral atom gains an e- is the ELECTRON AFFINITY

Electron Affinity Electron affinity decreases as you move down a group Increases as you move across a period Halogens have the highest electron affinities

Electron Affinity

Electronegativity (Pauling) Ability of an atom to attract (or remove) an e- from another atom Fluorine is most electronegative (4) Metals have Electronegativity of less than 2

Electronegativity Electronegativity decreases as you go down a group because the electrons are farther from the nucleus Electronegativity increases as you go to the right because atoms are more inclined to gain electrons in order to gain a full shell

Electronegativity

Summary of Periodic Trends Shielding is constant Atomic radius decreases Ionization energy increases Electronegativity increases Nuclear charge increases 1A Ionization energy decreases Electronegativity decreases Nuclear charge increases Atomic radius increases Shielding increases Ionic size increases 2A 3A 4A 5A 6A 7A • Elements with the highest ionization energies are those with the most negative electron affinities, which are located in the upper-right corner of the periodic table. • Elements with the lowest ionization energies are those with the least negative electron affinities and are located in the lower-left corner of the periodic table. • The tendency of an element to gain or lose electrons is important in determining its chemistry. • Various methods have been developed to describe this tendency quantitatively. • The most important method is called electronegativity (), defined as the relative ability of an atom to attract electrons to itself in a chemical compound. Rules for assigning oxidation states are based on the relative electronegativities of the elements — the more-electronegative element in a binary compound is assigned a negative oxidation state Electronegativity values used to predict bond energies, bond polarities, and the kinds of reactions that compounds undergo Trends in periodic properties: 1. Atomic radii decrease from lower left to upper right in the periodic table. 2. Ionization energies become more positive, electron affinities become more negative, and electronegativities increase from the lower left to the upper right. Ionic size (cations) Ionic size (anions) decreases decreases