E3 Describe the development of the modern periodic table E4 Draw conclusion about the similarities and trends in the properties of elements, with reference to the periodic table
By 1817 52 elements had been discovered. By 1863 62 discovered With 62 elements, scientists decided it was time to organize them. They started with the most obvious property atomic mass but soon realized this didn t work because of isotopes (the mass changed depending on which type of an element was present)
Dimitri Mendeleev was one of the first to publish a table. (1869) He used similar chemical properties as well as atomic mass to arrange the elements. (Ie some were solids, some very reactive, some high melting points, some very brittle, etc ) He noticed that when elements are listed according to their mass certain properties occur periodically (so in certain vertical families they all seem to have similar properties )
As more data became available we actually changed the table to increasing atomic number, which made the whole table make sense!!! (Isotopes no longer caused a problem to the pattern!) The PERIODIC LAW summarizes the organization of the periodic table: When elements are arranged in order of increasing atomic number, their properties are repeated periodically. Henry Mosely came up with the concept of atomic number.
Period- a horizontal row on the periodic table Sodium (Na) Family Name: Alkali metals Family #: 1 Group #: 3 How many elements in the 6 th period? 36
Families (a.k.a. Groups)- vertical columns on the periodic table in which elements have similar properties and chemical reactivity (because of similar electron configurations in the outermost energy levels (ie. Valence shells p-orbitals) Stability is achieved by having a full outer valence shell (full p or s-orbital), or even half full.
How many families on the periodic table? There are 18 families on the periodic table Interesting pattern: The last digit of the family # indicates how many valence e- there are in all elements in that family (works for Families 1-2, 13-18) How many e- would family 17 like to gain or lose? Family 2??
H = 1s 1 Li = [He]2s 1 Na = [Ne]3s 1 K = [Ar]4s 1 Rb = [Kr]5s 1 Cs = [Xe]6s 1 All elements in the same family end with the same valence configuration. (in this case 1s 1 ), same # of valence e-. Also, notice that the principle quantum number of the s orbital is the same as the period #
Because they have the same number of valence e-. And we know that it is the valence e- that are responsible for the way elements interact!
Metals (to the left of the staircase line ) Non-metals (to the right of the staircase line ) Metalloids/semi-metals/Semiconductors, along the staircase line (the line between metals and non-metals starts under Boron) Metalloids have properties of both metals and non-metals. The properties of elements change from metallic to non-metallic going from left to right across the table * Elements become more metallic going down a family in the periodic table
Recall: Properties of metals versus non-metals: Metals - Have a metallic shine/ luster - Usually solid at room temperature - Are malleable (don t shatter) - Are ductile (can be formed into sheets or wires) - Are good conductors of electricity Non-metals - Usually have a dull appearance/ lack luster - Usually gases at room temperature - Are brittle solids (i.e. are not malleable) - Are not ductile Do not conduct electricity
Increasing Metallic character Increasing Metallic character LEAST METALLIC MOST METALLIC
Include Li, Na, K, Rb, Cs, Fr most chemically reactive of the metals (must be stored under oil as they will react with air) react explosively with water
Includes Be, Ca, Mg, etc. very chemically reactive and are never found free in nature (same as alkali metals)
are highly reactive with metals, especially alkali metals highly toxic
Include He, Ne, Xe, etc EXTREMELY STABLE and only rarely react to form compounds (because of their full valence shell) All other elements want to be like these ones, and they will lose or gain electrons to become iso-electric (having some electron composition) as them. all are gases when given energy (with electricity or extreme heat) each gives off a unique color
1 2 3 4 5 6 7 Energy Levels So Li, Be, B, C, N, O, F, Ne are all in the same energy level, n=2, and as you move across from left to right, there are more e-)
Pretty Protons Excited Electrons
If there are more Pretty p+ in the nucleus, do you think they would have a greater pull on the exited e- who were all the same distance away? Heck yes! If there were a lot of e-(all at different distances) around the p+, would the outer e- be pulled in by the protons to the same extent as the ones that are closer? Out of sight out of mind??
Going across the row or period, the atomic size/radius decreases. Why do you think? This is because: e - and p+ are being added to the same energy level (same distance from the nucleus) which means more attractive forces between them (the p+ in the nucleus are pulling e- more strongly into the center because there are more of them and the e- are still at some distance away).
Pretty Protons Excited Electrons
Going down a group or family, the atomic size increases. Why? This is because: e - are being added to E.L. further away from the p+ in the nucleus. The p+ are not as close to all the e- anymore, and therefore don t exert as much pull on them. More e - means more e - repulsion between them. More e - shielding of the outer e - from the nucleus by the inner e -. Basically, you are adding more energy levels, and therefore the atom gets bigger!
The middle electrons (electrons in the energy levels between the nucleus and the valence electrons) are called the Shielding electrons, they "shield" the valence electrons from the force of attraction exerted by the positive charge of the p+ in the nucleus.
Excited Electrons are in each other s way (e- shielding= there s a lot of e- between the p+ and the outside e-. Likes repel, so there is more repulsion of e- than pulling by p+ outside e- are not being pulled in as much!! Pretty Protons
The smaller an atom is, the closer the p+ are to the valence e- and therefore, the e- are held more tightly by the force of attraction so it take a lot of energy to remove e- in smaller atoms. If the atom is larger, the nucleus is farther from the valence e-, and therefore they are held less tightly, so it is easier to remove e- from larger atoms.
Metallic ions, which are positive, lose electrons, so they are smaller than their atom forms. A) This is because the ions have more protons than electrons, which exert a stronger pull on the outer electrons. B) They might also lose an energy level (making it smaller)
For non-metallic ions, which are negative, the opposite is true(they gain more e-) The non-metallic ions are larger than their atoms because there are more electrons being added than protons so there is less attraction and more electron repulsion.
Going across a period, the ionization energy increases because: The atoms get smaller as you go across (attractive forces are greater within the atom), therefore, it s harder to pull the electrons away so more energy is needed to do that. Metallic nature of elements (Ability to Give up e-) decreases as you go across! Because the additional protons in the nucleus attract the e-, they cause the electrons to be less readily released.
Going down a group ionization energy decreases because: The atom gets bigger as you go down, so the electrons are further away from nucleus and not held tightly so they are easier to steal! Metallic nature increases! (e- are further away from the nucleus, and can be more easily removed) There is more electron repulsion between the e-, There is more electron shielding
Down a family more energy levels further from nucleus Electron shielding and repulsion!! Pretty Protons Low ionization energy won t take much to steal that valence e-!!!
In general ionization energy increases; as you go across a period, and decreases as you go down a family. There is a slight decrease in ionization energy for which a loss of an electron produces a full or half full orbital (because of their additional stability). Notice: The greatest ionization energies for a period are for the noble gases (they are VERY unreactive and do not want to lose an electron.
1 st Ionization energy means energy needed to remove 1 electron 2 nd Ionization energy means energy needed to remove a second electron 3 rd Ionization energy means energy needed to remove a third electron
Note: Which tend to attract/take electrons? Metals or Non-metals? NON-METALS!!! Elements get more non-metallic as you go across the table. As the atomic radius decreases, the protons pull electrons in more strongly (more + s pulling s) So as the radius gets smaller the tendency to pull in other e- increases. So going across the periodic table Electronegativity INCREASES!! (see the back of your periodic table)
Electronegativity (ability to attract e- to itself) Decreases as you move down the families, the atomic radius is increasing (atom getting bigger and not holding the e- so tightly), there is not a great tendency to pull in other e-.
Electronegativity
As you become more non-metallic your affinity for an e- increases. (non-metals are takers). If you are a metal you don t want extra e-, so your affinity is typically low. So in general, the affinity increases across a period, and decreases down a family.
Electronegativity
True or False Elements in the same family have similar properties? True! Elements in the same family will react in the same way? True! Why do they react the same way? Same # of valence e-!!
If Mg reacts with Cl - what will the formula be? MgCl 2 So then what will the formula be of Ca and Br (without doing any work?) CaCl 2 What about Ba and F? BaF 2 See the pattern!?
If Ca binds with P to form Ca 3 P 2 Predict the formula for Sr bonding with N? Sr 3 N 2 So the big idea is elements in the same family will make the same compounds because of their similar valence e-!!!