The Periodic Table Beyond protons, neutrons, and electrons

It wasn’t always like this…

Early PT Folks Johann Dobereiner Triads- groups of 3 with similarities/ trends C l, Br, I – the properties of Br were intermediate to those of C l and I Limited to some groups, not effective with others JAR Newlands (1864) Law of Octaves Every eight elements the pattern repeats itself, similar to a musical scale repeating every 8 notes Not generally well received; people thought him a fool

The Modern Periodic Table The original PT was arranged by mass By Dmitri Mendeleev and J Lothar Meyer in 1869 Mendeleev predicted the existence of unknown elements (which turned out to be Ge, Sc, and Ga), and predicted their properties from the patterns he saw Mendeleev corrected the assumed atomic masses for elements (In, Be, U)  These are reasons why he is credited with the first periodic table and is dubbed “The Father of the Modern Periodic Table” over Meyer

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Changes…. Henry Mosley changed the table to be organized by atomic number (Z) instead; it then more closely followed trends/ patterns

e - configuration and the PT PT also shows trends in electron configuration Groups are based upon electron configuration Alkali metals are #s 1 (# is period) Alkaline earth metals are #s 2 (# is period) Halogens #p 5 (# is period) Noble gases #p 6 (# is period) Transition metals d block (# is period -1) Inner transition metals are f block (# is period -2)

Blocks and l * * * orbital shape The blocks you already know correspond to the orbital of the last (outermost) e-, or valence e-s occupied

Patterns (Periods) and the PT We see patterns for many things, including Atomic number * (not a periodic pattern, but a pattern) Electron configuration Atomic radii Ionization energy Electron affinity Electronegativity Activity Density

The Periodic Law Mendeleev says "The properties of the elements are a periodic function of their atomic masses" We now say: “When atoms are arranged by increasing atomic number, the physical and chemical properties show a (repeating) pattern”

Periodic… Summed up: Properties of elements are periodic functions of their atomic numbers. Hence, we call the table of elements the PERIODIC table (go figure)

Octet Rule “Atoms gain, lose, or share electrons in order to create a full outer shell” This is typically going to be eight electrons H and He are exceptions; wanting to fill the 1s orbital  H may want to go to no electrons, which is considered “full” even though it is empty  H gains an electron to become H -, (same e configuration as He The law can be used to predict several properties

Nuclear Charge Nuclear charge – the attraction felt for an electron by the nucleus Electrons are both attracted to the nucleus and repelled by other electrons. The nuclear charge that an electron experiences depends on both factors. This effects all periodic properties

© 2009, Prentice-Hall, Inc. Effective Nuclear Charge The effective nuclear charge, Z eff, is found this way: Z eff = Z − S, Z = atomic number S = inner core e -

Atomic Radii Half the distance between adjacent nuclei ½ (2R)= atomic radius

Atomic Radii INCREASES as you go down a group This is because n increases DECREASES as you go across a period (Yes, this is counterintuitive) Due to the fact that you add e- as you add p+, so the nucleus is more positively charged, (electrons have the same negative charge) Results in each electron being more attracted to the (increasingly) more positive nucleus, and being pulled in closer Sort of like making a magnet more powerful- it will decrease the distance where it will pull objects towards it

Ionic Radii Cations (+) Smaller than the neutral atom The electrons have less repulsion, and pull in closer to the nucleus Anions (-) Larger than the neutral atom More electrons = more repulsion = larger electron cloud

Ionization Energy (Heretofore called IE) The amount of energy needed to remove an electron from an atom More specifically, an isolated atom of the element in the gas phase Measure in kJ/ mol A l (g)  A l (g) + + e - I 1 = 580 kJ/mol A l (g) +  A l (g) +2 + e - I 2 = 1815 kJ/mol

Why IE? Electrons want to hang around the atom (due to the protons in the nucleus pulling on them), so it takes energy to remove electrons In general The smaller that atom, the more energy it takes to remove an electron Because the electron is closer to the nucleus than in a larger atom The fewer electrons that atom possess, the harder it is to remove an electron Because it will hang on to them tighter as they are closer to the + charged nucleus; There is less repulsion between electrons too

IE, continued 1 st IE: energy needed to remove the first electron from an element 2 nd IE: energy needed to remove the second electron from an element

1 st IE

Successive IE There are also 3 rd, 4 th, 5 th, and so on IEs (which are successive IEs), until you can’t pull any more off It takes more energy to remove successive electrons than to remove the first Because…there are then more protons than electrons, and the stronger positive charge will then act on the remaining electrons to hold them to the atom (Remember that the charge on the nucleus increases while the charge on each electron remains the same, causing more pull by the nucleus on each individual electron)

Things to keep in mind… Remember (from coming up with the abbreviated electron configurations) that: Inner core electrons are those electrons from previous Noble Gas Valence electrons are the electrons that are on the exterior of an atom These are the electrons that are responsible for the behavior (properties) of the element

Successive IEs Are higher than the first Due to the fact that there is going to be more protons than electrons at that point, resulting in a stronger attraction on the remaining electrons than there was in the first place Basically increasingly larger jumps as each electron is removed One jump is usually much larger than the others, because once the inner core configuration is reached, electrons are removed from the inner core, taking a lot more energy

Successive IEs I1I1 I2I2 I3I3 I4I4 I5I5 I6I6 I7I7 Na Mg Al Si P Si Cl Ar

IE and the PT

Electron Affinity (EA) The energy change associated with the addition of an electron to a gaseous atom Negative values mean that energy is released when adding an e - more negative means more E released when adding an electron Wants an electron more than something with a more positive value Positive values mean that energy needs to be added to add an e - More positive means more E needed to add the electron Does not want an added electron; takes E to do it

The trend for EA is? EA becomes more positive moving down the PT EA becomes more negative from left to right Farther from the nucleus There are several exceptions to this The smaller the atom, the more e - -e - repulsion when adding electrons

EA trends

Electronegativity (Eneg) The ability of an atom to attract electrons in a bond Some atoms share electrons easily, others are electron hogs The ability to share is rated (usually) from 0 to 4 Elements with 0 Eneg share easily Elements with a high (close to 4) Eneg don’t share e - well

Electronegativity Trends If it normally goes +, it has a low Eneg If it normally goes -, it is has a high Eneg The smaller it is, the higher the Eneg The larger it is, the lower the Eneg Noble gases, which normally take no charge, we say have no Eneg values

Metallic character Metallic character is acting like a metal (conductive, shiny, malleable, etc.)) All elements possess from very low to very high metallic character. The scale is from Fr to F. Fr has the most metallic character and F has the least. In groups, metallic character increases with atomic number because each successive element gets closest to Fr. In periods, metallic character decreases when atomic number increases because each successive element gets closest to F.

Reactivity The nature (metal, non-metal, semi-metal) makes a difference in how an element’s chemical reactivity The trends are characterized by their nature

Metals reactivity trend In groups, reactivity of metals increases with atomic number because the ionization energy decreases. In periods, reactivity of metals decreases when atomic number increases because the ionization energy increases.

Nonmetals reactivity trend In groups, reactivity of non-metals decreases when atomic number increases because the electronegativity decreases Relate to size- it increases. In periods, reactivity of non-metals increases with atomic number because the electronegativity increases. Relate to size- radii decreases Remember, the radii would have an effect on this

Density: in general Density of solids is greatest Measured in g/cm 3 Highest in center of table (d- block) Density of gases Measured in g/L at Standard Temp &Pressure (STP, which is 1atm and 0°C) Increases as you go down a group Increases as you go across the table, then Decreases Density of liquids Measured in g/mL Density of Hg is greater than that of Br 2

Density

To sum it up…

Summary chart again