2 Development of Periodic Table Dmitri Mendeleev, Russian Chemist in the mid-1800’s arranged the 70 known elements into a systematic way.Put the names of elements on a card (Just like you arranged the pieces)He also put atomic masses, physical and chemical properties.When arranged in the order of atomic mass, he discovered a repeated, or periodic, pattern.He left blank spaces where he didn’t know the element for the spot.By doing this, he constructed the first periodic table for elements.
3 Mendeleev’s ProblemsWhy could most of the elements be arranged in the order of increasing atomic mass but a few could not?What was the reason for chemical periodicity?
4 Henry Moseley’s Solution Instead of arranging in atomic mass, in1913 Henry Moseley arranged in increasing order of nuclear charge. His work led to modern definition and recognition of atomic number.He discovers that it is the number of electrons that affect an atoms’ reactivity, or chemical property.
5 Parts of the Periodic Table Periods:The horizontal rowsPeriodic Law:When the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.Groups: (a.k.a. Families)The vertical columns
6 Parts of the Periodic Table, cont. Representative Elements:Group A elements. (The tall parts of the periodic table.)These can be divided into 3 groups:MetalsNonmetalsMetalloids
7 Metals Malleable-Very easy to be hammered or beaten into thin sheets Ductile-Very easy to produce a wireHigh luster-De-excitation of electrons causes shiny appearance of metal surfaceHigh electrical conductivity - Due to the free movement of electronsHigh thermal conductivityMost are solid at room temperatureMetallic Bond StrengthDirectly proportional to the heat of vaporization80% of all elements are metals
8 Nonmetals Upper right portion of the periodic table Generally nonlustrousGenerally poor conductors of electricitySome are gases at room temperature and some brittle solidsOne is even a liquid at room temperature
10 Metalloids These border the stair step, except for aluminum These have properties of both metals and nonmetalsSi, Ge, As, Sb, Te
11 Group 1 – Alkali MetalVery reactive!Not found in nature as free elementsCombine vigorously with most nonmetalsReact strongly with water to produce hydrogen gas and aqueous solutions of substances known as “alkalis.”Usually stored in kerosene (paraffin oil)
12 Group 2 – Alkali Earth Metal Harder, denser, and stronger than alkali metalHigher melting pointsNot so reactive compared to alkali metal; Still very reactive! Not found as free element
13 Group 3 – 12 Transition Metal Good conductors of electricity and have a high lusterLess reactive than alkali and alkali-earth metalSome (e.g. platinum and gold) are so unreactive that they do not form compounds easily.Some are found as free element.
14 Group 17 - Halogens Most reactive nonmetal React vigorously with most metals to form “salts”Need one electron to obtain stable noble-gas configurationF and Cl are gaseous.Br is liquid.Iodine is solid.
15 Group 18 – Noble Gases Lack chemical reactivity Very stable Octet electronic configuration in the outermost energy level (or “shell”)
19 Categories of electrons Core (inner) electrons – all those shared by the previous noble gasOuter electrons – those in the highest occupied energy levelSimilar chemical properties of elements in groups is a result of similar outer electron configurationsIn the main group, the group number equals the number of outer electronsValence electrons – those involved in bondingIn the main group, the outer electrons are valenceIn the transition metals can include some d electrons19
21 Metallic Behavior Metallic behavior increases left and down Metals tend to lose electrons in reactionsHighly metallic elements, are likely to make positive ionsLeast metallic elements are likely to make negative ionsIn middle are more likely to make covalent bonds21
22 Ion (Cation and Anion) # of electrons added/removed= # of charge Neutral atomAdd electronNegative ion (Anion)F+ 1e-F-1 (Anion)Neutral atomRemove electronPositive ion (Cation)Ca-2e-Ca+2 (Cation)# of electrons added/removed= # of charge
23 Main-group ions and the noble gas configurations Octet (Duplet) Rule2323
24 Cations and Anions Of Representative Elements +1+2+3-3-2-18.2
25 Periodic Law Be 1S2 2S2 2, 2 Mg 1S2 2S2 2p6 3S2 2, 8, 2 All the elements in a group have the same electron configuration in their outermost shellsExample: Group 2Be 1S2 2S22, 2Mg 1S2 2S2 2p6 3S22, ,Ca 1S2 2S2 2p6 3S2 3p6 4S22, , ,LecturePLUS Timberlake2525
26 Periodic Table and Electron Configuration Groups = s levelGroups 3-8 = p levelTransition = d levelLantanides = f levels1 s p1 p2 p3 p4 p5 p612d1 - d10456f1 - f1426
27 The relation between orbital filling and the periodic table Figure 8.13The relation between orbital filling and the periodic table27
30 Trends in Some Periodic Properties The physical and chemical behavior of the elements is based on the electron configurations of their atoms.e- configurations can be used to explain many of the repeating or “periodic” properties of the elements30
31 Defining metallic and covalent radii Figure 8.14 31
33 Definition of Some Periodic Properties Atomic Radius/Ionic Radius- size of atom/IonElectronegativity is a measure of an atom’s attraction for another atom’s electrons.Ionization energy: the energy required to remove an electron from an atom is ionization energy. (measured in kilojoules, kJ)Electron affinity is the energy change that occurs when an atom gains an electron (also measured in kJ).33
34 Atomic radii of the main-group and transition elements. VerticallyThe trend for atomic radius in a vertical column is to go from smaller at the top to larger at the bottom of the family.With each step down the family, we add an entirely new energy level, making the atoms larger with each step.HorizontallyEach step adds a proton and an electron (and 1 or 2 neutrons).Electrons are added to existing energy level.The effect is that the more positive nucleus has a greater pull on the electron cloud.The nucleus is more positive and the electron cloud is more negative.The increased attraction pulls the cloud in, making atoms smaller as we move from left to right across a period.34
35 Electron Shells and Sizes of Atoms Size in main group elements follow 2 general rules1. AR increases going down in a group2. AR decreases going left to right in a periodWhy?Moving down in a group:Zeff is constant because Z and S increase equallyElectrons in a higher n and therefore largerMoving across:Z increases and S does not (electrons in same shell don't shield each other well) so e- attracted more strongly and are closer/smallerIn transition metals there is an initial decrease in size but then remains relatively constant35
36 Periodicity of atomic radius Figure 8.16Periodicity of atomic radius36
37 Ranking Elements by Atomic Size SAMPLE PROBLEM 8.3Ranking Elements by Atomic SizePROBLEM:Using only the periodic table (not Figure 8.15), rank each set of main group elements in order of decreasing atomic size:(a) Ca, Mg, Sr(b) K, Ga, Ca(c) Br, Rb, Kr(d) Sr, Ca, RbPLAN:Elements in the same group increase in size and you go down; elements decrease in size as you go across a period.SOLUTION:(a) Sr > Ca > MgThese elements are in Group 2A(2).(b) K > Ca > GaThese elements are in Period 4.(c) Rb > Br > KrRb has a higher energy level and is far to the left. Br is to the left of Kr.(d) Rb > Sr > CaCa is one energy level smaller than Rb and Sr. Rb is to the left of Sr.37
38 Effective nuclear charge (Zeff) is the “positive charge” felt by an electron. Zeff = Z - s0 < s < Z (s = shielding constant)Zeff Z – number of inner or core electronsZeffCoreZRadius (pm)NaMgAlSi11121314101234186160143132
40 She’s unhappy and negative. IonsHere is a simple way to remember which is the cation and which the anion:This is Ann Ion.This is a cat-ion.She’s unhappy and negative.He’s a “plussy” cat!
41 Cation FormationEffective nuclear charge on remaining electrons increases.Na atom1 valence electronRemaining e- are pulled in closer to the nucleus. Ionic size decreases.11p+Valence e- lost in ion formationResult: a smaller sodium cation, Na+
42 Anion FormationA chloride ion is produced. It is larger than the original atom.Chlorine atom with 7 valence e-17p+One e- is added to the outer shell.Effective nuclear charge is reduced and the e- cloud expands.
43 Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed.
45 Ionic Size Cations are smaller than the parent atom Anions are larger than the parent atom.Ionic size increase down in a groupIn an isoelectronic series, the most negative ion is largest, the most positive is smallest.Atoms making more than one ion, most positive is smallest.4545
46 Ionization Energy This is the second important periodic trend. The atom has been “ionized” or charged.The number of protons and electrons is no longer equal.The energy required to remove an electron from an atom is ionization energy. (measured in kilojoules, kJ)The larger the atom is, the easier its electrons are to remove.Ionization energy and atomic radius are inversely proportional.Ionization energy is always endothermic, that is energy is added to the atom to remove the electron.
47 Ionization Energy (Potential) Draw arrows on your help sheet like this:
48 Ionization energy Trends: But Why? Generally as size decreases, IE1 increases.1. IE1 generally increases from left to right, some exceptions2. IE1 generally decreases going down in a group3. transition and f-block elements have much smaller variances in IE1But Why?Across - increasing Zeff and smaller size(e- closer to nucleus)exceptions - Be to B because s shields p and lowers ZeffN to O because repulsions in first paired electronDown - Zeff constant and larger so easier to ionize48
49 Periodicity of first ionization energy (IE1) Figure 8.17Periodicity of first ionization energy (IE1)49
50 First ionization energies of the main-group elements Figure 8.18First ionization energies of the main-group elements50
51 Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state.I1 + X (g) X+(g) + e-I1 first ionization energyI2 + X+(g) X2+(g) + e-I2 second ionization energyI3 + X2+(g) X3+(g) + e-I3 third ionization energyI1 < I2 < I3
52 Ranking Elements by First Ionization Energy SAMPLE PROBLEM 8.4Ranking Elements by First Ionization EnergyPROBLEM:Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE1:(a) Kr, He, Ar(b) Sb, Te, Sn(c) K, Ca, Rb(d) I, Xe, CsPLAN:IE decreases as you proceed down in a group; IE increases as you go across a period.SOLUTION:(a) He > Ar > KrGroup 8A(18) - IE decreases down a group.(b) Te > Sb > SnPeriod 5 elements - IE increases across a period.(c) Ca > K > RbCa is to the right of K; Rb is below K.(d) Xe > I > CsI is to the left of Xe; Cs is furtther to the left and down one period.52
53 Ionization energy Result - Rx only involve outer e- Second ionization energy IE2 – to remove 2nd e-Always larger than IE1as e- are removed, Z remains constant and remaining e- are harder to removeA huge increases occurs in ionization energies when core electrons are reached because of the lower shielding/higher Zeff.Result - Rx only involve outer e-53
54 The first three ionization energies of beryllium (in MJ/mol) Figure 8.19The first three ionization energies of beryllium (in MJ/mol)54
56 Identifying an Element from Successive Ionization Energies SAMPLE PROBLEM 8.5Identifying an Element from SuccessiveIonization EnergiesPROBLEM:Name the Period 3 element with the following ionization energies (in kJ/mol) and write its electron configuration:IE1IE2IE3IE4IE5IE61012190329104956627822,230PLAN:Look for a large increase in energy which indicates that all of the valence electrons have been removed.SOLUTION:The largest increase occurs after IE5, that is, after the 5th valence electron has been removed. Five electrons would mean that the valence configuration is 3s23p3 and the element must be phosphorous, P (Z = 15).The complete electron configuration is 1s22s22p63s23p3.56
57 ElectronegativityElectronegativity is a measure of an atom’s attraction for another atom’s electrons.It is an arbitrary scale that ranges from 0 to 4.The units of electronegativity are Paulings.Generally, metals are electron givers and have low electronegativities.Nonmetals are are electron takers and have high electronegativities.What about the noble gases?
58 ElectronegativityYour help sheet should look like this:
59 Overall ReactivityThis ties all the previous trends together in one package.However, we must treat metals and nonmetals separately.The most reactive metals are the largest since they are the best electron givers.The most reactive nonmetals are the smallest ones, the best electron takers.
60 Overall ReactivityYour help sheet will look like this:
61 Electron Affinity What does the word ‘affinity’ mean? Electron affinity is the energy change that occurs when an atom gains an electron (also measured in kJ).Where ionization energy is always endothermic, electron affinity is usually exothermic, but not always.
62 Electron AffinityElectron affinity is exothermic if there is an empty or partially empty orbital for an electron to occupy.If there are no empty spaces, a new orbital or PEL must be created, making the process endothermic.This is true for the alkaline earth metals and the noble gases.
63 Electron AffinityYour help sheet should look like this:++
64 Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion.X (g) + e X-(g)F (g) + e X-(g)DH = -328 kJ/molEA = +328 kJ/molO (g) + e O-(g)DH = -141 kJ/molEA = +141 kJ/mol
65 electron affinitiesenergy required to put an e- on an atom (making it negative or more negative)EA1 usually negativeEA2 always positiveIrregular trend: increases going right and upRemember: Full shells best, full subshells good, and half full is kinda coolCl and other halogens need only 1 e- so most exothermic reactionnoble gases have endo because new electron in higher nMg, Be endo because new e in higher lN endo because new e paired65
66 Electron affinities of the main-group elements Figure 8.20Electron affinities of the main-group elements66
67 Trends in three atomic properties Figure 8.21Trends in three atomic properties67