Presentation on theme: "5-3 Electron Configurations and Periodic Properties"— Presentation transcript:
1 5-3 Electron Configurations and Periodic Properties
2 PreClass: What does the numerical value of the period (ie PreClass: What does the numerical value of the period (ie. Period 4) represent in addition to just physically locating a particular row on the Periodic Table?The period = n value of the outermost electronsIe. If n=4, than all elements on the 4th period have their outermost electrons in the 4 energy level !!
3 Beaker Breaker1. Mendeleev’s periodic table had the elements ordered in increasing ________2. What is the element in the same Period that follows potassium, K ?3. What is the name of the group on the far left of the Periodic table?4. What element ends with 3p1 ?
4 Atomic RadiusOne-half the distance between the nuclei of identical atoms that are bonded together
5 Group Trend for Atomic Radii Down a group the radius of an atom gets largerThe principal quantum number increases and the cloud grows by one shell
6 Period Trend for Atomic Radii Across a period the radius of an atom gets smallerYou’re adding electrons to approximately the SAME region (same “n”) while the electrons are being pulled in more tightly as nuclear charge increases (more protons)Exception: Noble Gas Family – atoms don’t interact and pull together like in other atoms because atoms already have full outer level
7 QuestionsOf the elements Li, O, C, and F, identify the one with the largest atomic radius and the one with the smallest atomic radiusOf the elements Br, At, F, I and Cl, identify the one with the largest atomic radius and the one with the smallest atomic radius
8 AnswersLi = largest F=smallestF = smallest At = largest
9 Why does the atomic radii not decrease very much as you move across the d-sublevel? Increasing # inner electrons shield outer level electrons from nucleus
10 Why would the atomic radius of hafnium (#72) be LESS than that of zirconium (#40)? Hf has such a greater nuclear charge
11 Beaker BreakerWhich has the largest and which has the smallest atomic radius of the following:C, Ge, Sn, Si2. Which has the largest and which has the smallest atomic radius of the following:K, Cu, As, Br
12 Ionization energy, IEEnergy required to remove an electron from a neutral atom (or first ionization energy, IE1)Ion: charged particleIonization: process of an electron being lost or gained from an atom which results in the formation of an ion (Na+ and Cl-)
13 Group Trend for Ionization Energies Down a group the ionization energy of an atom generally gets smallerElectrons are at a greater distance from the nucleusOuter electrons are shielded from the nucleus by inner electrons
14 Period Trend for Ionization Energies Across a period the ionization energy increasesNuclear charge gets greater while atomic size decreases
15 Examples Which of the following has the largest ionization energy? Na, Mg, P, ClWhich of the following has the smallest ionization energy?Be, Mg, Ca ,Sr
17 Beaker BreakerWhich alkaline earth metal has the lowest ionization energy?Which halogen has the largest atomic radius?
18 What will determine whether an electron is easily lost or not What will determine whether an electron is easily lost or not? (4 things)
19 4 Factors Affecting Ionization Energy Nuclear charge (greater the charge, the greater the IE)Shielding Effect (greater the shielding, the lower the IE)Radius (the greater the radius, the less the IE)Sublevel configuration (an electron from a half-full or full sublevel requires more IE)
20 2nd Ionization energy, IE2 is the energy required to take a 2nd electron away from an atom. Why is IE2 always greater than IE1? Why is the IE2 of Na so much greater that the greater of Mg ?? (Hint: look at the electron config.)
21 Multiple Ionization Energies Energy required to remove the 2nd, 3rd, etc. electron from an atomIE3 > IE2 > IE1 because remaining electrons will be held more tightly as the electron repulsion decreases and the cloud is pulled in more tightly
22 Why does the IE go UP as you go down a d-sublevel group but go down in an s-sublevel? “f” sublevel has minimal shielding effect while the nuclear charge continues to grow (“Lanthanide Contraction”)
23 Electron Affinity p. 147Energy change that occurs when an electron is acquired by a neutral atomMost atoms RELEASE energy when they acquire an electron: A + e- A- + energyenergy has negative signSome atoms gain energy when they acquire an electron: A + e- + energy A-energy has positive sign; atom is unstable & loses electron spontaneously
24 Group Trend for Electron Affinities Down a group the electron affinity of an atom tends to get smallerAlthough there is an increasing nuclear charge, there are more levels so the size is greaterAdding an energy level usually dominates!
25 Period Trend for Electron Affinity Across the p-sublevel, the energy change increases (becomes more negative)Electron config is close to being full and the size is smaller
26 Why is the electron affinity of nitrogen so low when compared to carbon or oxygen? Adding an electron to carbon half fills the 2p sublevelAdding an electron to nitrogen forces the config to go from stable (half-filled) to less stable (no spec arrangement)
27 Multiple Electron Affinities It is always more difficult to add a 2nd electron to an already negatively charged ion..therefore, all 2nd electron affinities are positive
28 Beaker BreakerWhich alkaline earth metal has the largest electron affinity?Which of the following has the smallest electron affinity Na, Mg, P or Cl
29 Ionic Radii of CationsCation: positive ionFormed by an atom losing electron(s)Always smaller because electron cloud is smaller (less repulsion) & sometimes even one less energy level!
30 Ionic Radii of Anions Anion: negative ion Formed by an atom gaining electron(s)Always larger because electron cloud is greater (more repulsion among electrons)
31 Group Trend for Ionic Radii Down a group the ionic radius of an atom generally gets largerElectrons are at a greater distance from the nucleus (higher E level) and have more shielding
32 Period Trend for Ionic Radii Metals (left side):form cationsCationic Radius: Decreasing ionic radius as nuclear charge increases without adding an energy levelNon-metals: form anionsAnionic Radius:Decreasing ionic radius as nuclear charge increases without adding an energy level
35 Valence ElectronsElectrons available to be lost, gained, or shared in the formation of chemical compoundsOften the outermost electrons because they are held most loosely
36 What would be the # of valence electrons in…….. CalciumLithiumChlorinecarbon
37 What would be the # of valence electrons in…….. Calcium – 2 : 4s2Lithium – 1: 2s1Chlorine – 7: 3s23p5Carbon – 4: 2s22p2
38 Beaker Breaker How many valence electrons does Te have? Which has a smaller ionic radius Na+1 or Ca+1 ?
39 ElectronegativityMeasure of the ability of an atom in a chemical compound to attract electronsFluorine is the MOST electronegative element – assigned an arbitrary value of 4.0All other values are relative to F3 highest values: F – O - N
40 Group Trend for Electronegativity Tend to decrease down a group (or stay the same) as the atoms gets larger
41 Period Trend for Electronegativity Tend to increase across the period as the atoms gets smaller, the nuclear charge becomes greater, and the atom is getting closer to a noble gas configuration
42 Determine the likely charge for the following elements: Ca, O, Al Write the noble gas configuration of the elementDetermine if electrons will be LOST or GAINED to make the element stableID the noble gas whose electron configuration by losing/gaining these electronsWrite the formula for the ionID it as a cation OR anion
43 Determine the likely charge for the following elements: Ca, O, Al Ca: [Ar]4s2Ca will LOSE 2 e-Ca now has the Ar configCa+2cationO: [He]2s22p4O will GAIN 2 e-O now has a Ne configO-2 , an anion
44 Al Al: [Ne]3s23p1 Al will LOSE 3 e- Al now has a Ne config Al+3 , a cation
45 How do d-block elements form ions? Electrons in the highest occupied sublevel are always removed first