The Periodic Table Beyond protons, neutrons, and electrons
It wasn t always like this
Early PT Folks n Johann Dobereiner n Triads- groups of 3 with similarities/ trends n Cl, Br, I the properties of Br were intermediate to those of Cl and I n Limited to some groups, not effective with others n JAR Newlands (1864) Law of Octaves n Every eight elements the pattern repeats itself, similar to a musical scale repeating every 8 notes n Not generally well received; people thought him a fool
The Modern Periodic Table n The original PT was arranged by mass n By Dmitri Mendeleev and J Lothar Meyer in 1869 n 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 n 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. n Henry Mosley changed the table to be organized by atomic number (Z) instead; it then more closely followed trends/ patterns
Patterns (Periods) and the PT n We see patterns for many things, including n Atomic number *(not a periodic pattern, but a pattern) n Electron configuration n Atomic radii n Ionization energy n Electron affinity n Electronegativity n Activity n Density
The Periodic Law n Mendeleev says "The properties of the elements are a periodic function of their atomic masses" n We now say: When atoms are arranged by increasing atomic number, the physical and chemical properties show a (repeating) pattern n Hence, we call the table of elements the PERIODIC table (go figure)
The Modern Periodic Table n Consists of boxes containing: n Element name and symbol n Atomic number and atomic mass n Groups: n Columns (up and down) n 18 of them n Periods: n Rows (left to right) n 7 of them n Representative elements: n elements in groups 1,2 and 13-18 n Have a wide range of chemical and physical properties
Metals n Conductors of heat and electricity n Make cations (lose e - to become + charged) n Malleable (made into sheets) n Ductile (made into wire)
Nonmetals n Are a brittle solid or a gas n Make anions (gain e - to become - charged) n Covalently bond to each other
Semi-metals (AKA Metalloids) Characteristics of both metals and nonmetals More metallic as you go down PT
Valence electrons and the PT PT also shows trends in valence electrons. Elements in: n Column 1 have 1 valence e- n Column 2 have 2 valence e- s n Columns 13 to 18 have valence e- s equaling the column # - 10 n Column 13 have 3 valence e- s (13-10) n Column 14 have 4 valence e- s (14-10) n Column 15 have 5 valence e- s (15-10)
Periods, e - configuration and the PT n Period trends are seen in the electron configuration n For columns 1,2 and 13-18, the period on the table matches the energy level (ring, n value) n Alkali metals are #s 1 (# is period) n Alkaline earth metals are #s 2 (# is period) n Halogens #p 5 (# is period) n Noble gases #p 6 (# is period) n Different trend for d and f blocks n Transition metals #d (# is period -1) n Inner transition metals are #f (# is period -2)
e - configuration and the PT n Group trends are also seen in the electron configurations n Columns 1 and 2 end in s (s 1 and s 2 ) n Columns 13 to 18 end in p (p 1 to p 6 ) n Transition metals end in d (d 1 to d 10 ) n Inner transition metals end in f (f 1 to f 14 )
Blocks and * orbital shape l* * The blocks you already know correspond to the orbital of the last (outermost) e-, or valence e-s occupied
Octet Rule n Atoms gain, lose, or share electrons in order to create a full outer shell n This is typically going to be eight electrons n n H and He are exceptions; wanting to fill the 1s orbital H gains an electron to become H -, with the same electron configuration as He H may want to go to no electrons, which is considered full even though it is empty H + and He +2 would have no electrons left n The law can be used to predict several properties
+1 +2 Common Ions of Elements Variable, always + +3 +/-4-3 -2-1 0
Effective Nuclear Charge n Nuclear charge the attraction felt for an electron by the nucleus n Electrons are both n attracted to the nucleus and n repelled by other electrons. n The nuclear charge that an electron experiences depends on both factors. n This effects all periodic properties
Atomic Radii Half the distance between adjacent nuclei ½ (2R)= atomic radius
Atomic Radii n The radius increases as you go down a group n This is because n increases n The radius decreases as you go across a period (Yes, this is counterintuitive) n Due to the fact that you add e- as you add p+, so the nucleus is more positively charged, and each electron has the same negative charge n n 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 n Cations (+) n Smaller than the neutral atom n n Anions (-) The electrons have less repulsion, and pull in closer to the nucleus n Larger than the neutral atom n More electrons = more repulsion = larger electron cloud
Ionization Energy (Heretofore called IE) n IE is the amount of energy needed to remove an electron from an atom n (specifically, an isolated atom of the element in the gas phase) n Measure in kj/ mol Al (g) àal (g)+ + e - I 1 = 580 kj/mol Al (g)+ à Al (g) +2 + e - I 2 = 1815 kj/mol
IE, continued n The Energy needed to remove the first electron from an element is the 1 st IE n The Energy needed to remove the second electron is known as the 2 nd IE
Successive IE n 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 n It takes more energy to remove successive electrons than to remove the first n Due to the fact that 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 n (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)
Why IE? n Since electrons (-) want to hang around the atom (due to the + protons in the nucleus pulling on them), it takes energy to remove electrons n In general n The smaller that atom, the more energy it takes to remove an electron n Because the electron is closer to the nucleus than in a larger atom n The fewer electrons that atom possess, the harder it is to remove an electron n Because it will hang on to them tighter as they are closer to the + charged nucleus; n also, the less repulsion between electrons
1 st IE
Things to keep in mind n Remember (from coming up with the abbreviated electron configurations) that: n Inner core electrons are those electrons from previous Noble Gas n Valence electrons are the electrons that are on the exterior of an atom n These are the electrons that are responsible for the behavior (properties) of the element
Successive IEs n Are higher than the first n 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 n Basically increasingly larger jumps as each electron is removed n 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 n Much bigger difference between positive nucleus and negative electron
Successive IEs I 1 I 2 I 3 I 4 I 5 I 6 I 7 Na 495 4560 Mg 735 1445 7730 Al 580 1815 2740 11600 Si 780 1575 3220 4350 16100 P 1060 1890 2905 4950 6270 21200 Si 1005 2260 3375 4565 6950 8490 27000 Cl 1255 2295 3850 5160 6560 9360 11000 Ar 1527 2665 3945 5770 7230 8780 12000
Electronegativity (Eneg) n The ability of an atom to attract electrons in a bond n Some atoms share electrons easily, others are electron hogs n The ability to share is rated (usually) from 0 to 4 n Elements with 0 Eneg share easily n Elements with a high (close to 4) Eneg don t share e - well
Electronegativity Trends n If it normally goes +, it has a low Eneg n If it normally goes -, it is has a high Eneg n The smaller it is, the higher the Eneg n The larger it is, the lower the Eneg n Noble gases, which normally take no charge, we say have no Eneg values
Electronegativity Trends
Metallic character n n n n Metallic character is acting like a metal (conductive, shiny, malleable,etc) All elements possess from very low to very high metallic character. n The scale is from Fr to F. n 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 n The nature (metal, non-metal, semi-metal) makes a difference in how an element s chemical reactivity n The trends are characterized by their nature
Metals reactivity trend n In groups, reactivity of metals increases with atomic number because the ionization energy decreases. n In periods, reactivity of metals decreases when atomic number increases because the ionization energy increases.
Nonmetals reactivity trend n In groups, reactivity of non-metals decreases when atomic number increases n because the electronegativity decreases n Relate to size- it increases. n In periods, reactivity of non-metals increases with atomic number n because the electronegativity increases. n Relate to size- radii decreases n Remember, the radii would have an effect on this
Density: in general n Density of solids is greatest n Measured in g/cm 3 n Highest in center of table (d- block) n Density of gases n Measured in g/l at Standard Temp &Pressure (STP, which is 1atm and 0 C) n Increases as you go down a group n Decreases as you go across the table n Density of liquids n Measured in g/ml n Density of Hg is greater than that of Br 2
Density
Density v atomic number
Summing it up (again)
Summary chart again