The Periodic Table. Beyond protons, neutrons, and electrons

Transcription

1 The Periodic Table Beyond protons, neutrons, and electrons

2 It wasn t always like this

3 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

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5 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

6 Ekasilicon

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

8 e - configuration and the PT n PT also shows trends in electron configuration n Groups are based upon electron configuration 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 Transition metals d block (# is period -1) n Inner transition metals are f block (# is period -2)

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

10 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

11 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

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

13 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 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 n 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

14 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

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

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

18 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 Results in each electron being more attracted to the (increasingly) more positive nucleus, and being pulled in closer n Sort of like making a magnet more powerful- it will decrease the distance where it will pull objects towards it

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20 Radii (2) * The radii for the Noble gases is INCORRECT- they should be smaller!

21 Ionic Radii n Cations (+) n Smaller than the neutral atom n Anions (-) n 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

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23 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) à Al (g) + + e - Al (g)+ à Al (g) +2 + e - I 1 = 580 kj/mol I 2 = 1815 kj/mol

24 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

25 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)

26 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

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28 1 st IE

29 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

30 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

31 Successive IEs I 1 I 2 I 3 I 4 I 5 I 6 I 7 Na Mg Al Si P Si Cl Ar

32 IE and the PT

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

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

35 EA trends

36 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

37 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

38 Electronegativity Trends

39 Metallic character n Metallic character is acting like a metal (conductive, shiny, malleable,etc) n 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. n In groups, metallic character increases with atomic number because each successive element gets closest to Fr. n In periods, metallic character decreases when atomic number increases because each successive element gets closest to F.

40 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

41 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.

42 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

43 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

44 Density

45 Density v atomic number

46 Summing it up (again)

47 Summary chart again