Natural Sciences I. Lecture 15: Elements and the Periodic Table

Transcription

1 Natural Sciences I Lecture 5: Elements and the Periodic Table Classifying MATTER Chapter 0 of your text opens with a discussion of the types of chemical bonds, and also includes a review (pp. 4-4) of the properties of solids, liquids and gases. We talked about the states of matter in lecture, and we'll deal with bonding after we gain a little more background in how matter is classified... EXAMPLES Pure Substances Compounds water salt sugar alcohol MIXTURES Mixtures show variability in chemical composition, both within a single sample and from one sample to the next. We recognize two types of mixtures... Homogeneous mixtures: A good example is a solution, where one component is dissolved in another to varying degrees. The word "solution" is essentially synonymous with "homogeneous mixture". Examples include: sea water Elements hydrogen oxygen carbon iron etc. Matter Homogeneous Mixtures salt water metal alloys etc. All matter is made of chemical elements usually but Mixtures not always in combination... Heterogeneous Mixtures sand dirt rocks etc. stainless steel (an alloy of iron with other metals such as nickel or chromium) H O Na + Ca ++ Cl Mg ++ CO 3

2 Heterogeneous mixtures: These are aggregates with physically distinct and separable parts; sometimes also called "mechanical" mixtures. Examples include: microscopic beach sand view quartz grain soil concrete most rocks PURE SUBSTANCES shell fragment iron oxide Pure substances are materials that are the same throughout and have a fixed composition (one sample is the same as the next); again, there are two types... Compounds A compound is a pure substance that can be decomposed by chemical change (e.g., heating, burning) into simpler substances in a fixed ratio. Example: Calcium carbonate (CaCO 3). A major constituent of skeletal material, shells, etc.; active ingredient in TUMS and other antacids. CaCO 3 CaO (calcium oxide) CO The calcium oxide and carbon dioxide are always produced in the same mass ratio. Elements A pure substance that cannot be broken down by any chemical or physical means is called an element. We now understand, of course, that each element corresponds to a particular "kind" of atom, but getting to this point required quite a lot of detective work by early chemists. Your text describes some of the highlights and abandoned theories of this processes.

3 3 One of the key figures in the development of our modern view of matter was John Dalton. In 0 he published his New System of Chemical Philosophy, which stated five principles:. Matter is made up of individual atoms that cannot be subdivided (we now understand that this is not strictly true today). Each chemical element is made up of a particular kind of atom 3. Atoms are unchangeable (again, this is understood today to be the case only as long as nuclear processes are not involved) 4. Chemical elements can combine with one another to make compounds 5. Chemical reactions can reorganize atoms into different compounds, but the number of atoms of each element does not change in the process (this is sort of a corollary to no. 3). The history of the discovery of the elements... number of known elements electrolysis (Volta's battery) air liquified spectroscope developed Lavoisier's work with gases (see p. 4) synthetic elements year

4 4 Lavoisier used reactions with gases to deduce the existence of many elements. His 9 list of 33 known elements was correct in 3 cases ( were actually compounds; he also considered light and heat to be elements). Many of his studies involved heating materials in the presence of gases in such a way that he could deduce the amount of gas consumed and the mass ratio of solid or liquid to gas... glass flask O O O O water level rises if oxygen is consumed O O O trapped air (oxygen) O water mercury converted to red oxide by reaction with oxygen in known proportion Following Lavoisier, several major advances in technology took place in the 9th century such as the battery, the spectroscope, and cooling systems capable of liquifying gases. These made it possible to deduce the existence and atomic numbers of many more elements. By 940, the list of naturally-occurring elements was complete as we now know it. These range from hydrogen (Z=) to uranium (Z=9), with some omissions and complications: The heaviest elements thorium and uranium are radioactive; bismuth is the heaviest stable element (Z=3). technetium (Z=43) and promethium (Z=6) do not occur naturally plutonium (Z=94; heavier than uranium) exists in minute quantities in nature In the years following the World War II, a number of radioactive elements of higher atomic number than uranium were produced artificially by bombardment with subatomic particles in particle accelerators. We won't worry about these until we discuss radioactivity (soon)...

5 5 A digression on Natural Abundances... As noted on page 4, the naturally-occuring elements span a range in atomic number from to 9. The abundances of these elements vary enormously both in the Earth and in the cosmos, and from one planetary body to another even within our own solar system. Only a handful of elements are abundant in the Earth... CRUST: the outer "skin" 5-50 kilometers thick MANTLE: ~ half Earth's radius; mostly Mg, Si, O CORE: mostly iron; some nickel Earth's Crust Whole Earth element symbol wt % element symbol wt% oxygen silicon aluminum iron calcium sodium potassium magnesium O Si Al Fe Ca Na K Mg Many of these values are somewhat "model-dependent" iron oxygen silicon magnesium nickel calcium sulfur aluminum chromium phosphorus cobalt Fe O Si Mg Ni Ca S Al Cr P Co The abundances of the remaining 0-or-so elements are so low that we call most of them "trace elements"

6 6 Atomic Masses, Isotopes John Dalton used the fixed mass ratios of combining elements to deduce their relative masses. He knew, for example, that one gram of hydrogen combines with grams of oxygen to make 9 grams of water. He assumed that each "particle" (that is, each molecule) of water contained atom each of hydrogen and oxygen. He could not know that each "particle" of water actually contains two hydrogen atoms, so he was wrong in this case about their mass ratios. Dalton's water "particle" H O g hydrogen + g oxygen = 9 g water deduced oxygen:hydrogen mass ratio = : (the measurement is correct, but the deduction is wrong) actual water molecule H H O unless you know that the molecule actually contains 3 atoms, it is not possible to deduce the mass ratio...so Dalton made lots of mistakes! Chemists improved on Dalton's approach in the century following his work, but it wasn't until the early 0th century that atomic and molecular masses could be accurately measured. The mass spectrometer was developed for this purpose by F.W. Aston (a former assistant to J.J. Thomson the fellow who had measured the charge/mass ratio of the electron a few years earlier). (next page)

7 An early MASS SPECTROGRAPH accelerating grid magnet evacuated "flight tube" ( ) detector ( ) sample inlet ("source") (+) electron beam to ionize sample magnet magnetic field bends ion beam (depends on mass) Example: A mass spectrum for chlorine gas (Z=) signal intensity atomic mass (u) Note: The atomic mass (u) is the sum of the number of protons and the number of neutrons. With the development of mass spectrometry, chemists began to realize that even a single chemical element often had atoms of more than one mass. These atoms behave the same way chemically because they have the same number of protons and electrons, but they differ in their number of neutrons. The are called ISOTOPES of the element as in the case of chlorine illustrated here.

8 Recognizing hydrogen as the lightest (lowest-z) element, Dalton had originally assigned a mass of unit to this element, and normalized all others to it. Today, the accepted convention is to assign a mass of.00 to carbon (6 protons + 6 neutrons) and normalize all other elements to that. Note, however, that isotopes of other elements do not have integer mass values for example, an oxygen atom containing protons and neutrons does not have a normalized mass of exactly 6 (but it is very close at ). The difference arise because the differing binding energies of the nuclei. Chlorine has two stable isotopes: chlorine-35 chlorine-3 35 ( Cl; mass units; 5.53% of all Cl atoms) 3 ( Cl; mass units; 4.4% of all Cl atoms) The atomic mass of chlorine is determined by the weighted average of the masses of these two isotopes: average atomic mass = ( amu) (5.53%) + ( amu) (4.4%) = 6.4 amu amu = amu...so the atomic mass of chlorine is reported (on the periodic table, for example) as Chlorine summary atomic number symbol isotopes (%) atomic "weight" Cl chlorine-35 (5.53) chlorine-3 (4.4) amu

9 9 Partial summary of other (stable) isotopes isotope mass abundance (%) atomic "weight" H H Be N 5 N 6 O O O F Ne Ne 0 0Ne 3Al

10 0 Patterns of element behavior As the number of known elements proliferated during the 9th century, chemists began to discern systematic trends in chemical and physical properties of the elements as a function of atomic weight. One of the earliest.. ideas was that of the triad, put forth by German chemist Johann Dobereiner in 9. He grouped elements in threes on the basis of similar properties, and noted some interesting trends within each group... "triads" element lithium (Li) sodium (Na) potassium (K) calcium (Ca) strontium (Sr) barium (Ba) atomic wt. (amu) average atomic wt density 3 (g/cm ) boiling pt. O C chlorine (Cl) bromine (Br) iodine (I) In Dobereiner's time, there existed no explanations for the existence of these systematic trends. Decades later (69), Russian chemist Dmitri Mendeleev and German physicist Lothar Meyer made things even more interesting by publishing detailed The average atomic weight of the first and third elements is very close to that of the middle element (this is true to some extent of other properties such as density, melting point, and boiling point) classifications that provided a glimpse into the true extent of the systematic trends. Both charts were based upon atomic weights. Being a chemist, Mendeleev also considered chemical properties; Meyer, being a physicist, incorporated physical properties into his scheme. Mendeleev is probably more famous today...

11 Systematic behavior of boiling points... boiling point (C) Li C H He Na Si Sc V Co Mn Ge K Zn Rb Ru Cd Sn La Cs Ne Ar Kr Xe atomic number Note that every eighth element is a noble gas He, Ne, Ar, Kr and Xe. Mendeleev did not know about atomic numbers, and he did not have all the data shown in this plot, but he was aware of the "octave" systematics these determined the appearance of his periodic table... Here's part of a periodic table arranged according to Mendeleev's rules (his first table of 69 actually contained 6 elements): oxide group period R O R O 3 R O 5 R O RO RO RO 3 I II III IV V VI VII VIII H He 3 Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Ti V Cr Mn Fe Co Ni Cu Zn As Se Br Kr Mendeleev was the first to recognize the existence of gaps in his table (see next page)...

12 In the years following 69, Mendeleev published many iterations of his original classification, which eventually evolved into the modern-day periodic table. He was so confident about his approach that he predicted not only the existence of elements unknown at the time, but also their chemical properties. Here is an example section of one of Mendeleev's tables where he left spaces for elements he knew would eventually be discovered... Sc Ga Mg Al Si P Ca Ti V Zn As Sr Y Zr Nb Ge Note: This is a piece of Mendeleev's chart, so it's not quite the same as the one you know Mendeleev's periodic table was unsatisfying in a couple of ways. Remember that a principal basis of his classification was the atomic weight, and it did not escape the notice of his critics that the progression in atomic weights was not perfectly regular. More importantly, perhaps, is that there was no explanation for the dependence of chemical and physical properties upon atomic weight. As physicists working in the early 0th century began to decipher the structure of the atom (that is, when protons and neutrons were discovered), it eventually became clear that the observed properties of elements depend upon the atomic number, not the atomic weight. This leads to the modern periodic law When the elements are listed in order of atomic number, similar physical and chemical properties periodically recur Today, this law is embodied in the PERIODIC TABLE.

13 3 periods H He Li Be B C N O F Ne 3 NaMg Al Si P S Cl Ar 4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn GaGeAs Se Br Kr 5 Rb Sr Y Zr NbMo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 6 Cs Ba La Hf Ta W ReOs Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Ce Pr NdPmSm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf IA H Rb Cs Fr groups or families IVA VIA VIIIA IIA IIIA VA VIIA He IVB VIB IIIB VB VIIB VIIIB IB IIB 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 Sr Ba Y La Ra Ac Zr Hf Nb Ta Mo W Tc Re Ru Rh Pd Os Ir Pt Ag Au Cd Hg In Tl Sn Pb Sb Bi Te Po I At Xe Rn Elements belonging to the same group or family have similar properties (this is more true of some groups than of others)

14 4 Most of you probably remember that the chemical properties of the elements (which affects their placement in the periodic table) are actually determined by the electronic structure of the atoms in particular, the filling or "completeness" of the outermost orbitals or energy levels. Recall the sequence of filling orbitals as we discussed in lecture 0... energy 6s 5s 4s 3s s s 5p 4p 3p p 4d 3d 4f Your text provides several tables and other aids to help you build the electronic structure of progressively higher-z atoms. Here are a couple of approaches: group VIIIA element helium neon argon krypton 'xenon radon Z K L M N 3 O P Q These "shell" designations correspond to principle quantum nos. - group IIA group IA lithium sodium potassium rubidium cesium francium beryllium magnesium calcium strontium barium radium The similar "completeness" of the outermost (highest energy) orbitals imparts similar chemical properties to these groups of elements group VIIA fluorine chlorine bromine iodine astatine

15 5 A color-coded table to help you understand the connection between chemical properties and the filling of electron "shells"... period H He IVA VIA VIIIA IA IIA IIIA VA VIIA Li Be B C N O F Ne 3 4 IVB VIB Na Mg Al Si P S Cl Ar IIIB VB VIIB VIIIB IB IIB 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 55 Ba 56 La 5 Hf Ta W Re Os Ir Pt Au Hg Tl 9 0 Pb Bi 3 Po At 4 5 Rn 6 Fr Ra Ac One of the most important aspects to appreciate is that at the first orbitals filled s (n = period) as you move element of each period from left to d (n = period - ) (a group IA element) you right across p (n = period) begin filling a new main the periodic level (n) by adding the table f (n = period - ) first electron to an s sublevel. Another key to understanding the diagram is that each period contains a number of elements corresponding to the number of electrons required to fill certain sublevels s, p, d, f.

16 6 Atoms participate in chemical reactions and compound formation through processes involving their outer (highest energy) electrons. These may be lost, gained, or shared with others atoms. The chemical reactivity of elements and their tendency to combine with certain other elements in specific ways depends upon the upon the extent to which the outer electron energy levels (orbitals) are filled. Unreactive elements: Noble gases element helium (He) neon (Ne) argon (Ar) krypton (Kr) xenon (Xe) radon (Rn) electron configuration s 6 [He]s p 6 [Ne]3s 3p 0 6 [Ar]4s 3d 4p 0 6 [Kr]5s 4d 5p [Xe]6s 4f 5d 6p The noble (inert) gases are "noble" in the respect that they don't ordinarily participate in reactions with other elements. This is because their outer electron orbitals are complete. Reactive elements: alkali metals element lithium (Li) sodium (Na) potassium (K) rubidium (Rb) cesium (Cs) francium (Fr) electron configuration [He]s [Ne]3s [Ar]4s [Kr]5s [Xe]6s [Rn]s The alkali metals are extremely reactive because they have an unpaired s electron in the outermost orbital. This electron is easily lost to create a cation with a + valence and the electron structure of a noble gas. Alkali metals react violently with water to make an alkaline solution; they do not exist in the metallic state in nature. In the chemical sense, an element is considered a metal if it tends to lose electrons to form a positive ion. Elements with, or 3 outer electrons tend to do this. Na Na + e

17 Reactive elements: halogens element fluorine (F) chlorine (Cl) bromine (Br) iodine (I) electron configuration 5 [He]s p 5 [Ne]3s 3p 0 [Ar]4s 3d 4p [Kr]5s 4d 5p The halogens are extremely reactive because they are one electron short of completing their outermost (p) orbitals. Acquisition of one additional electron yields an anion with a valence of - and the electron structure of a noble gas. An element is considered a nonmetal if it tends to gain electrons to form a negative ion. Elements with 5, 6 or outer electrons tend to do this F e F Electron dot notation is one way in which we show pictorially the outer electron structure of elements and their tendency to gain or lose electrons... H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Ga Ge As Se Br Kr Rb Sr In Sn Sb Te I Xe Cs Ba Tl Pb Bi Po At Rn Fr Ra

18 One more way of classifying the elements period IA H He IIA IIIA VA VIIA Li Be B C N O F Ne 3 4 IVB VIB Na Mg Al Si P S Cl Ar IIIB VB VIIB VIIIB IB IIB 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 55 Ba 56 La 5 Hf Ta W Re Os Ir Pt Au Hg Tl 9 0 Pb Bi 3 Po At 4 5 Rn 6 Fr Ra Ac 9 Summary of element types metals nonmetals semiconductors With reference to the periodic table, we recognize four basic types of elements. These are: Noble Gases at the end of each period in family VIIIA unreactive colorless gases at STP Representative Elements include both metals and non-metals in the A families the number of valence electrons (outermost shell) is the same as the group number IVA Transition Elements all are metals in the B families filling of d or f orbitals d and s (or f and s) electrons involved in reactions VIA VIIIA