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CHAPTER 7- PERIODIC PROPERTIES OF THE ELEMENTS • Periodic table- One of the most significant tools chemists have for organizing and remembering chemical facts • In Chapter 6- what did we say was primarily responsible for the properties/reactivities of elements? • Electron configuration- repeating pattern on periodic table • Responsible for periodic nature of the elements • Elements in same column (group)= same number of valence electrons • What were valence electrons again? • Elements in same group- what else is special about them? • Why do elements in the same group have similar properties? • What you haven’t heard- elements in same group share properties, but also have differences • Ex: Oxygen vs. Sulfur- both in group 6A- bond/react similarly • Oxygen: colorless gas at room temp • Sulfur: yellow solid at room temp • These differences can also be attributed to electron configuration • What shell (value of n) are the valence electrons of O in? • What shell are the valence electrons of S in? SEC. 7.1 • Mendeleev- created one of the first widely accepted periodic table • Inspired by his favorite card game, solitaire, Mendeleev created a set of cards of the elements- one card for each element currently discovered. • Each element card had information on it: the element’s name, atomic mass, properties, and compounds they formed with other elements • Mendeleev organized his periodic table in order of increasing atomic mass- he spread all the cards out, placed them in a big line according to increasing mass • When ordered this way, he noticed a pattern arise among the elements- similar properties actually repeated themselves in a regular pattern (a periodic pattern) • Periodic Table Build! • You are given a set of element cards- all info on these cards should be used in order to organize the elements you are given into the first periodic table (correctly) • Discovery of chemical elements has been an ongoing process since ancient times • Though many elements are stable- tend to form compounds • Not many elements found in elemental form naturally- why’s that? • Many elements went undiscovered due to the fact that they naturally form compounds • Early 1800s- advances in chem allow for isolation of elements from their compounds • These advances= many elemental discoveries • With number of elements increasing, the demand for a method/way of organizing them increased • Many scientists tried many different classifying schemes • In steps Mendeleev with his crazy table, rambling on about how his way of organizing the elements was the best way • How did he organize it? • Why did you end up leaving holes in your periodic table? • What did you end up doing with the holes that were left? • Mendeleev ended up with the same holes in his periodic tablehe insisted that elements with similar properties go in the same families- this forced him to leave holes in his table • Using his awesome table, Mendeleev boldly went were no man had gone before and decided he would predict the existence of these undiscovered elements- he went as far as predicting their atomic mass, melting point, color, density, specific heat, boiling point, and the compounds it would form with oxygen and chlorine • He called the element hole under Aluminum “eka- Aluminum” and the element hole under Silicon “eka- Silicon” • When Gallium and Germanium were discovered, they very closely matched the properties Mendeleev has predicted • This proved how useful Mendeleev’s periodic table was- it can be used to predict properties of undiscovered elements Property Mendeleev’s prediction for eka-Silicon (1871) Germanium properties (1886) Atomic weight 72 72.59 Density 5.5 g/mL 5.35 g/mL Specific heat 0.305 (J/gK) 0.309 (J/gK) Melting Point High 947 °C Color Dark gray Grayish white Oxide Formula XO2 GeO2 Chloride Formula XCl4 GeCl4 Boiling Point of chloride compound A little under 100 °C 84 °C • Even though Mendeleev’s table was the bees knees, it had its problems • Do you remember which elements broke the trend of increasing atomic mass? Why did you place them out of order? • After Rutherford’s model of the atom, a physicist named Henry Mosely decided he would shoot high energy electrons at atoms • These atoms emitted X-rays of specific and unique frequencies- frequency emitted generally increased as atomic mass increased • Assigned each element a specific whole number values based on the frequency of X-ray emitted- atomic number • Stated atomic number= number of protons in nucleus • Ordered periodic table in order of increasing atomic number- this is how the modern periodic table is arranged • Ordering in increasing atomic number- clarified the problems seen in Mendeleev’s table • This new arrangement also allowed for the identification of holes/undiscovered elements- property predictions as well- led to discovery of even more elements SEC. 7.2- EFFECTIVE NUCLEAR CHARGE • Electrons- negative charge • Nuclei- contain protons- positive charge • Electron attracted to nucleus- atom properties depend on electron configurations and how strongly their outer electrons are attracted to the nucleus • Coulomb’s Law= strength of interaction between 2 electrical charges depends on magnitudes of charge and distance between them • Attractive force between electrons and nucleus increases as nuclear charge increases • Decreases as electrons move further away • Many electron atoms- all electrons attracted to nucleus and repelled by one another at the same time • Due to number of repulsions, it is difficult to measure all the forces • However, we can estimate the net attraction of each electron to the nucleus- must consider how single electron interacts with “average” environment created by nucleus and other electrons • This allows us to treat each electron as if it were moving in a net electric field created by nucleus and electron density of other electrons • Effective Nuclear charge (Zeff)- the nuclear charge felt by an electrontakes into consideration the “shielding” effects the inner shell electrons have; this is the “net” nuclear charge felt by individual electron • Shielding- caused by the repulsions between electrons • Effective nuclear charge will always be smaller than the actual nuclear charge • Valence electron- attracted to nucleus in atom, but also repelled by inner shell electrons- the repulsions from inner shell electron partially “cancel out” some of the attraction felt from the nucleus (shielding) • Zeff = Z – S • Zeff= effective nuclear charge, Z= nuclear charge (# protons in nucleus), S= shielding constant • S is portion of nuclear charge that is “cancelled out” by repulsion of inner shell electrons • S is normally close to the value of inner shell electrons • Electrons in same valence shell are not good at shielding one another ZEFF FOR A SODIUM ATOM • Periodic trend: Zeff for valence electrons increases as we move from left to right across a period • Why do you think this is? What changes as we move from left to right? • Going down a group, Zeff slightly increases, but not by much- not nearly as much as it does moving across a period • Would expect the Zeff to be about the same for elements in same group- why? SEC. 7.3- SIZES OF ATOMS AND IONS • Size of atoms and ions is a somewhat strange concept, but is also extremely important • Atoms- not hard, spherical objects- if you remember from chapter 6, electron clouds are not really “hard” boundaries- no sharp cut off for the boundary at which electron distribution becomes zero • We define atomic size in several ways- mostly based on the distance between atoms in various situations • Argon atoms- when these gas particle collide, they easily bounce off one another without reacting/interacting all that much • This is due to the fact that the electron clouds of these gas atoms will not mesh together to any significant extent – no attraction because these atoms are not charged, and the clouds will actually somewhat repel one another • When atoms collide in this way, the collision is sometimes referred to as a “non-bonding” collision- since the atoms are simply ricocheting off one another • Nonbonding atomic radius- distance between two nuclei during a “non-bonding” collision • When atoms are bonded to one another- like Cl2- there is some sort of attractive force between the bonded atoms- hence why they are bonded • Due to this attractive force between bonded atoms, the nuclei of the bonded atoms will be closer to one another than they would be if they were not bonded- the bond kind of squeezes their electron clouds closer together • Bonding atomic radius- the distance between the nuclei of bonded atoms • How does bonding atomic radius compare to nonbonding atomic radius? • Through observation,/experiments, scientists have gathered data on the bonding atomic radius for every element- each element assigned its own bonding radius value • With these values, we are able to tell how long a bond will be between any combination of bonded atoms in molecules • Ex: For C, bonding atomic radius= 0.77 Å, and Cl= 0.99 Å • In the compound CCl4, we can estimate the bond length of a C-Cl bond to be 1.76 Å (In reality, the bond length is 1.77 Å, which is pretty close to the estimated value) PERIODIC TRENDS IN ATOMIC RADII • Two trends for atomic radius: • 1- Atomic radius tend to increase from top to bottom (going down) a group. Why do you think this is? What happens to the value of n as we go down a group? • 2- Atomic radius tends to decrease when moving from left to right across a period. Major influence for this is the Zeff. Why does Zeff have an effect on atomic radius? PERIODIC TRENDS IN IONIC RADII • Ionic radii- determined by distance between nuclei of ions • Size of ion depends on effective nuclear charge, number of electrons it possesses, and the orbitals in which its valence electrons are located • When a cation forms (positive ion), does this mean that electrons were gained or lost? • When an anion forms (negative ion), does this mean that electrons were gained or lost? • If cations form when electrons are lost, how do you think the ionic radius of a cation compares to that of the regular atomic radius? • If anions form when electrons are gained, how does the ionic radius of an anion compare to that of the regular atomic radius? • Ionic radius of cation < atomic radius • Ionic radius of anion > atomic radius • Sooo… cations are smaller than their “parent” atoms, and anions are larger than their “parent” atoms • Also, for ions with same charge, size increases as we move down a group • Based on periodic trends for atomic radius and ionic radius, order to following atoms and ions in order of increasing size (smallest to largest): Mg+2, Ca+2, and Ca • Isoelectronic series: groups of ions all containing same number of electrons • Ex: O2-, F-, Na+, Mg2+, Al3+ … all have how many electrons • If members of isoelectronic series are listed in order of increasing atomic number, effective nuclear charge will increase. Why? • If nuclear charge increases, what will happen to the ionic radius? • List the ions above in order of increasing ionic radius SEC. 7.4- IONIZATION ENERGY • How easily electrons can be removed from atoms/ions has major impacts on chemical behavior • Ionization energy- minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion • First ionization energy (I1)- energy required to remove the first electron from a neutral atom • Second Ionization energy (I2)- energy needed to remove the second electron • This pattern continues for third, fourth, fifth ionization energies (continues until atom is out of electrons) • Which ionization energy do you think is the smallest? Why? • First ionization energy for the sodium atom is the energy required for the process: • I2 for sodium is energy associated with this process: • Ionization energies- always positive, due to the fact that energy must be applied to the atom to pull an electron away • First ionization energy is always the smallest- the first electron is easiest to pull away- partially due to the fact that the atom is neutrally charged at this point • Higher ionization energies= more difficult to remove electrons • This is due to the fact that after the first ionization, the atom becomes positively charged- creating a stronger attraction between the electrons and the nucleus • There is also a trend when it comes to inner shell vs. valence shell electrons- which electrons do you think would be easiest to remove? • There is a sharp increase in ionization energy once all valence electrons have been removed- greater increase than from the first to second ionization energy (if more than 1 valence electron) • Inner shell electrons- much more difficult to remove than valence electrons- why? • The fact that valence electrons have such low ionization energies than the inner shell electrons supports the fact that valence electrons are mostly responsible for how something bonds/ chemically reacts • The valence electrons will be the only ones willing to transfer/share- form ionic/covalent bonds • Inner electrons too tightly bound to nucleus to be lost from atom or even be shared with another atom • Predict which atom will have the largest second ionization energy: Calcium, Sodium, and Sulfur PERIODIC TREND OF FIRST IONIZATION ENERGY • First ionization energy, I1, typically increases going left to right across the same period (some exceptions to this rule) • Ionization energy typically decreases going down a groupas atomic number increases • The s and p block element show a larger range in values of I1 than the d block (transition metals) do. Generally, the ionization energy of the transition metals increases slowly as we move from the left to the right across a given period. • In general, smaller atoms have higher ionization energies- why do you think this happens? • The energy needed to remove an electron depends on: • Effective nuclear charge • Distance of the electron from the nucleus • Increasing effective nuclear charge or decreasing the distance between the electron and nucleus will increase the attraction between the electron and the nucleus • As attraction to nucleus increases, it becomes more difficult to remove the electron- more attraction = more required energy= higher ionization energy • Moving left to right across a period, what happens to: • Effective nuclear charge? • Atomic size? • So energy needed to remove electron will: • When moving down a group, what happens to: • Effective nuclear charge? • Atomic size? • So, energy needed to pull electron away will: • Arrange the following from smallest to largest first ionization energy: Ne, Na, P, Ar, K SEC. 7.5- ELECTRON AFFINITIES • Electron affinity- energy change that occurs when an electron is added to a gaseous atom • Measures attraction (affinity) of an atom to an added electron • Measures the ease in which an atom will gain an electron • Electron affinities are typically negative energy values (this indicates that energy is released when electron added) • If the electron affinity value is positive, this means the atom will not gain an electron • The more negative the electron affinity is, the more the atom will want to gain an electron • No real periodic trend when it come to electron affinitieshowever, some groups of elements have characteristically low/high affinities • The Halogens- Have the most negative electron affinity values… why? • The Noble Gases- do you think their electron affinities will be negative or positive? Why? SEC. 7.6- METALS, NONMETALS, AND METALLOIDS • Elements can be broadly grouped into the categories of metals, nonmetals, and metalloids • Metalloid staircase- marked on your periodic tables • Left of staircase- metals (any exceptions?) • Right of staircase- Nonmetals • Touching staircase on two sides- metalloids (Boron, silicon, germanium, arsenic, antimony, tellurium)- form a border between metals and nonmetals • What category do most elements on the periodic table belong to? • Metallic character- a measure of how much an element exhibits the physical and chemical properties of metals • Metallic character- Increases moving down a group; decreases moving left to right across a period METALS • Shiny luster; various colors- often silvery • Good conductors of heat and electricity (this is due to their metallic bonding) • Most are Malleable- can be pounded into thin sheets • Most are Ductile- can be pulled into a thin wire • High melting points- all are solid at room temp except one… which one? • Tend to have low ionization energies- what does this mean they tend to do with their valence electrons? • Tend to form what charges? • When metals undergo chemical changes, they typically oxidize (Oxidation is Loss) due to low ionization energies/ electron loss • For s block metals- they will lose electrons in the outermost s shell- forming positive charges- either 1+ or 2+ charges form • For p block metals- They will either lose just the outer p electrons, or both the outer s and p electrons • Al3+ - what block did the electrons this metal lost belong to? • Sn2+ - what block did the electrons this metal lost belong to? • Sn4+- what block did the electrons this metal lost belong to? • Charges formed by transition metals do not follow a distinctive pattern- often able to form more than one positive ion • When metals bond with nonmetals, they typically form an ionic compound • When metals left outside, typically combine with the nonmetal oxygen (abundant in our air) to form an ionic compound: • Ni + O2 → • Which one of these reactants- Ni or O2, will be easier to pull electrons away from? How do you know this? • Ionic bonding- one atom will lose electrons- the other will gain- which element is losing electrons in this reaction? Which one is gaining? • Whenever metals react with oxygen, we call the products metal oxides- metals oxides are fairly common ionic compounds due to the abundance of oxygen in nature • Metal oxides have a special characteristic- most of them are basic when in solution • When metal oxides are able to dissolve in water (if soluble) they react to form metal hydroxides: • Metal oxide + water → metal hydroxide • If you look at the net ionic equation, the basicity of metal oxides is due to the oxide ion, which reacts with water as follows: • Metal oxides also display their basicity when reacting with acids: • Metal oxide + acid → salt + water • What kind of reaction does that look like? NONMETALS • Nonmetals- vary greatly in their physical appearance • No luster, and are poor conductors of heat and electricity • Melting points- typically lower than those of metals (Notice that all gases on the periodic table are nonmetals) • Not malleable/ductile- may be very hard, brittle, or soft • Nonmetals- how do their ionization energies compare to those of metals? What about their electron affinities? • So when nonmetals bond with metals, what will they typically do? • Ex: Aluminum reacting with bromine • When nonmetals bond ionically (with other metals) they will gain enough electrons to fill their outermost occupied p subshell • When nonmetals bond with other nonmetals- this forms a covalent bond (electrons are shared) • Nonmetal oxides- acidic! • When dissolved in water, nonmetal oxides react to form an acid • Nonmetal oxide + water → acid • Ex: • Nonmetal oxides will also react with bases to form salt and water • Nonmetal oxide + base → salt + water • Ex: METALLOIDS • Metalloids- kind of like the strange child of metal/nonmetal • Metalloids have properties that are both metallic and nonmetallic- depends on which metalloid you are dealing with • Ex: Silicon- looks like a metal, but when you hit it with a hammer, it shatters (not malleable like metal) • Several metalloids- most notably silicon- are semiconductors and are used in computer chips/circuits • Pure Silicon= electrical insulator… but if you add specific impurities, its electrical conductivity dramatically increases- this makes it possible to control electrical conductivity by simply controlling chemical composition SEC. 7.7- GROUP TRENDS FOR THE ACTIVE METALS • We have seen that elements in the same group posses general similarities • There are trends found without groups- using periodic table and our knowledge of electron configuration to examine the chemistry of the alkali metals, and the alkaline earth metals GROUP 1A- ALKALI METALS • Alkali Metals= metals in group 1A • Soft, metallic solids, silvery, metallic luster, and good conductors of heat/electricity • They have a fairly low density and melting point (compared to other metals) • Moving down the alkali metals: melting point decreases, density increases, what about atomic radius? First ionization energy? • For each row (period) on the periodic table, the alkali metals have the lowest I1 value- what does this reflect about its outer s electron? • Due to the low ionization energy- alkali metals lose electrons very easily= very very very reactive • All will form what charge? • Alkali metals only exist as compounds- partially due to high reactivity • Will combine directly with most nonmetals • Alkali metals will combine with hydrogen to form hydrides and sulfur to produce sulfides: Na + H2 → Na + S → • These reactions happen with all alkali metals • When reacting with H2, alkali metals will give up their electron to H, which will then form a hydride ion: • All alkali metals are violently reactive with water • When they combine with water, they produce hydrogen gas and a metal hydroxide • The reaction between alkali metals and water is extremely exothermic- which means it releases a lot of heat- most of the time this is enough heat to ignite the H2 gas that is also produced… this makes a boom • How do you think the violence of this reaction changes as you go down a group? • Alkali metal ions= colorless in compounds/ solutions- But, when placed in flame, they all emit a characteristic color of light • When placed in flame, the alkali metal ions are reduced to gaseous metal atoms, and the valence electron absorbs energy from the heat of the flame • What happens when the single valence electron of an alkali metal absorbs energy? • When electrons become excited, they will want to immediately return to ground state • Ground state? • What will the excited valence electron have to do in order to return to ground state? • When releasing energy, the excited electron releases a photon with unique energy- due to this unique energy, the photon emitted will have a special wavelength- this will correspond with a specific color of light • Every metal ion releases a unique color of light when placed in a flame- why? • Ex: Sodium will release yellow light after being excited- valence electron in 3s jumps to 3p when excited- when returning back to ground state (3p to 3s) it releases a photon with a wavelength of 589 nm (yellow light) GROUP 2A: THE ALKALINE EARTH METALS • Alkaline earth metals are harder than alkali metals, have higher melting points, and are more dense • Going down the alkaline earth metals group, melting points decrease, density increases, atomic radius? First ionization energy? • How do you think their reactivity compares to the alkali metals? • What about reactivity going down the alkaline metal group? • The way the alkaline earth metals react with water shows the trend of reactivity (increases going down the alkaline earth metals)- none are as reactive as the alkali metals • Beryllium does not react with water, but every other alkaline earth metal does • Ex: Magnesium and water (all other alkaline earths- besides beryllium- react with water the same way) • Which electrons will the alkaline earth metals lose? • What ion will all alkaline earth metals form? • Alkaline earth metal with chlorine and oxygen: • Heavier alkaline earths (lower in group) are even more reactive to nonmetals • Alkaline earth metals are similar to alkali metals when in the presence of a flame • When heated to very high temperatures, their valence electrons will become excited- once excited they will release photons with a characteristic color (energy/wavelength) to return to ground state • Flame test analysis can help us identify the type od metal in an unknown compound SEC. 7.8- GROUP TRENDS OF SELECTED NONMETALS HYDROGEN • Very strange- doesn’t know exactly what to do with its valence electronwill be happy losing it, or gaining 1 more • Hydrogen does not lose its valence electron as easily as the alkali metals • When bonding with other nonmetals, hydrogen will share its valence electron instead of losing it • When reacting with alkali metals, will gain an electron to form H(hydride ion) • In aqueous chemistry, hydrogen loses its electron to form H+, hydrogen cation (acid!) GROUP 6A- THE OXYGEN FAMILY • Not many shared physical characteristics- mix of nonmetals, one metalloid, and a metal • Oxygen- gas at room temp, the rest of the group are solids • Oxygen, sulfur, selenium- typical nonmetals; telluriummetalloid; Polonium- radioactive metal • All have 2 valence electrons and will typically form what type of charge: • Oxygen- found in 2 naturally occurring molecular forms • O2= dioxide (typically referred to as just “oxygen”), and O3 = ozone • O2 and O3 are allotropes- different forms of the same element in the same state (O2 and O3 are gases) • O2 is necessary for life- we use it for respiration • O3 is found in upper atmosphere and polluted air- damaging to health, but also necessary for life on Earth- ozone absorbs a large part of the UV radiation emitted from the Sun • O2 has a great tendency to attract electrons (oxidizes other elements)- when combining with metal, forms oxide ion O2• Sulfur also has a tendency to gain electrons from other elements to form sulfides containing the S2- ion (also does this in combination with metals) GROUP 7A- THE HALOGENS • Group 7A= the Halogens (group beginning with fluorine) • All are typical nonmetals (astatine is rare and radioactive, so not much data) • All (except astatine) are diatomic • Going down the halogen group- what do you think happens to melting point? • All have different colored vapors- F2= pale yellow, Cl2= yellow-green, Br2= reddish-brown, I2= violet vapor • Highly negative electron affinities- meaning what? • The halogens are the most reactive nonmetals due to their extremely negative electron affinities • What ion will all halogens form? • Fluorine- most reactive, removes electrons from almost any substance it comes into contact with • Halogens will react with metals to form metal halides: • Halogens also react with hydrogen to form gaseous hydrogen halide compounds: GROUP 8A- THE NOBLE GASES • Group 8A= The Noble Gases- far right of the periodic table • All are nonmetals and gases at room temperature • All are monatomic- single atom, not a molecule (not diatomic) • Completely filled s and p subshells- this makes them happy (full valence shells) • All have very large ionization energies- meaning what? • All have positive electron affinities- meaning what? • Stable electron configuration= very unreactive • Do not readily form compounds • Radon- radioactive