Periodic Properties of the Elements Chapter 8

Periodic Properties -- a property that is predictable based on the element’s position within the periodic table Atomic Radius (size) Ionization Energy (ionization potential) Electron Affinity Electronegativity Metallic Character

The Periodic Table The modern periodic table was developed by Dmitri Mendeleev – Based on periodic law Which states that when elements are arranged in order of increasing mass, their properties recur periodically Mass increased from left to right… elements with similar properties fell into the same columns

Mendeleev's arrangement allowed him to predict the existence of elements not yet discovered and even predicted (accurately) their properties… Why?? How ???

What determines properties? Electron Configurations – show the particular orbitals that are occupied for an atom These are (by default) in the ground state – But can be used to show electrons “jumping” to an excited state

The e- configuration tells us hydrogen’s one electron is in the 1s orbital (the lowest orbital first – ground state) – The highest energy level in the entire configurations = valence electrons So how many valence electrons does Hydrogen have??

Orbital Diagrams The electron configuration can be represented slightly differently by using orbital diagrams The electron spin is represented by the direction of the arrow (up or down)

Electron Spin Two fundamental aspects of electron spin: 1.Spin, like a (-) charge, is a basic property of all atoms. One electron does not spin more or less than another (all have the same amount) 2.Only two possibilities: +1/2 (up) and -1/2 (down) 4 th and final QN  Spin quantum number (m s )

Pauli Exclusion Principle The Pauli Exclusion Principle states no two electrons in an atom can have the same four quantum numbers – Even if they have the same n, l, m l … then they must have different m s (+1/2 and -1/2)

Degenerate Orbitals Orbitals that are present but do not have any electrons in them are said to be degenerate – Like the 3s, 3p, and 3d orbitals in hydrogen – H 1 : 1s 1 – There are no electrons past the 1s orbital

Energy Associated In general, the lower the l quantum number, the lower the associated energy E (s orbital) < E (p orbital) < E (d orbital) < E (f orbital)

Shielding Electrons feel the repulsive effect from other electrons AS WELL AS the attractive effect from the positively charged nucleus The electrons between cause a screen effect or “shielding” that decreases the attractive force of the (+) nucleus to outer electrons

Effective Nuclear Charge (Z eff ) The third electron of lithium experiences Z eff The effective nuclear charge for this third electron is approximately from the nucleus and 2- from two shielding electrons, leaving 1+ (the charge)

Penetration As the electron “penetrates” the cloud, it feels the effective of the nucleus more fully because it is less shielded If the electron could somehow get closer to the nucleus it would feel the full 3+ attraction The s orbitals penetrate deeper than the p orbitals and thus are lower in energy (closer to the nucleus, smaller r (radius) – s < p < d < f

Things to notice… In the 4 th and 5 th principal levels, the effects of penetration become so important that the 4s orbital lies lower in energy than the 3 d orbitals and the 5s lies lower in energy than the 4 d orbitals

Things to notice… The energy separations between one set of orbitals and the next become smaller – The relative energy ordering of these orbitals can actually vary among elements – These variations result in abnormalities in e- configurations of the transition metals and their ions

“Conceptual Connection” Lets take a look at page 325

Electron Configurations A systematic way of illustrating the increasing energy of electrons – They fill the lowest energy first and work their way up (if ground state, of course) – Two electrons per orbital  opposite spins m s QN of +1/2 and -1/2

Principles / Rules This patter of orbital filling is known as the Aufbau Principle (German for “building up”) Hund’s Rule states when filling degenerate orbitals, electrons fill them singly first, with parallel spins, then double up with opposite spins

Hund’s Rule = Pizza Party

e - configs… a summary Electrons occupy orbitals so as to minimize the energy of the atom.. Lower energy orbitals fill before moving on to the higher energy orbitals Orbitals hold no more than two electrons with opposite spins (Pauli exclusion principle) – no 2 electrons can have the same 4 QN’s) Electrons fill orbitals singly with parallel spins before doubling up (Hund’s Rule)

Review The s sublevel has only one orbital and can only hold 2 electrons The p sublevel has three orbitals and can hold 6 electrons The d sublevel has five orbitals and can hold 10 electrons The f sublevel has seven orbitals and can hold 14 electrons

Practice Write out the electron configurations for each of the following elements: (a)Mg (b) P(c) Br(d) Al What element has the following configuration: 1s 2 2s 2 2p 6 3s 2 3p 5 4s 2 3d 4

More practice Write out the orbital diagram for sulfur and determine the number of unpaired electrons

More practice Write out the electron configuration for Ge and identify the valence electrons and the core electrons Try the same for phosphorus

Using the “f” orbitals

Valence Electrons The chemical properties are largely due to the number of valence electrons This explains the reactivity of hydrogen and the inertness of helium

The Octet Rule Noble gases are least reactive Elements with configurations similar to noble gases are most reactive The formation of cation and anions to obtain their octet…

Atomic Size Most of the atom’s size is from the electrons The orbitals are just “high probabilities” How can we define atomic size?

Radius Two Ways: – Van der Waals radius – one-half the distance between adjacent nuclei in an atomic solid (meaning they are not bonded but just ‘frozen’ up against one another) – Bonding atomic radius or covalent radius one-half the distance between two bonded atoms As is diatomics or other BONDED molecules

Atomic Radius **TREND #1** A more general term, the atomic radius, refers to the average bonding radii determined from many bonded and non-bonded measurements

Atomic Radius As you move down a column (or family) in the periodic table, atomic radius increases – Why?? As you move to the right across a period (or row) in the periodic table the atomic radius does what??? – It decreases!! Think about Z eff and shielding…

A General Trend: Atomic Radius

Think about it… Why does it take 1312 kj/mol of energy to remove the 1s electron from a hydrogen atom but 5251 kJ/mol to remove it from He + H1s 1 He + 1s 1

The core electrons efficiently shield electrons in the outermost principal energy level from nuclear charge, but outermost electrons doe not efficiently shield one another from nuclear charge

Transition Metals These do not change all that much As another proton is added, as is another electron – This goes into the highest energy level (outside) Because the 4s orbital fills before the 3d, this ratio stays roughly constant, which makes the radius stay roughly constant

Practice Choose the larger atom from each and explain your choice… trends are not reasons… (a)N or F(b) C or Ge(c) N or Al(d) Al or Ge **Page 338 for additional reasoning

Ions When making cations – subtract electrons (Li) 1s 2 2s 1 (Li + ) 1s 2 2s 0 When making anions – add electrons (F) 1s 2 2s 2 2P 5 (F + ) 1s 2 2s 2 2P 6

Transition Metal Ions Transitions are a little more complicated… – Remove electrons in the highest n-value orbitals first, even if this does not correspond to the reverse order of filling (V) [Ar] 4s 2 3d 3 (V 2+ ) [Ar] 4s 0 3d 3

Explanations Paramagnetic – contains unpaired electrons and therefore is attracted by external magnetic fields Ag [Kr] 5s 1 4d 10 Diamagnetic – contains only paired electrons and therefore slightly repelled by external magnetic fields (preferred) Zn 2+ [Ar] 4s 0 3d 10

Practice Write the e- configs and orbital diagrams for each of the following ions. State whether they are diamagnetic or paramagnetic (a)Al 3+ (b) S 2- (c) Fe 3+ Diamagnetic DiamagneticParamagnetic Hint: write out the atom’s config first and then do what needs to be done to make the ion

Ionic Radii **Trend #2** The radius of an ion is the ionic radius (same concept as atomic radius) What happens to the radius as of an atom when it becomes a cation? An anion?

Ionic Radii What happens to the radius as of an atom when it becomes a cation? An anion?

Ionic Radii In general, cations are much smaller than their corresponding atoms Anions are much larger than their corresponding atoms

Practice Choose the larger atom or ion from each (a)S or S 2- (b) Ca or Ca 2+ (c) Br - of Kr S 2- is larger Ca is larger???????? The Br - is larger than the Kr atom because Br - has one fewer proton than Kr which results in a lesser pull on the electrons (Z eff ) and thus a larger radius.

Isotopes Think about isotopes… Would you expect C-12 and C-13 to have different atomic radii?? The isotopes of atoms have the same atomic radii because neutrons are relatively small and are neutral in charge.

Ionization Energy (IE) ** Trend #3** The ionization energy (IE) of an atom or ion is the energy required to remove an electron from the atom or ion in the gaseous state

Ionization Energy (IE) The energy required to remove the first electron is called the first ionization energy The second  the second ionization energy

Ionization Energy (IE) Na (g)  Na + (g) + 1e - IE 1 = 496 kJ/mol Na + (g)  Na 2+ (g) + 1e - IE 2 = 4,560 kJ/mol Much harder because it does not want to lose that second electron

Summary of IE IE usually decreases as you move down a column because they are farther away from the nucleus and are held less tightly IE generally increase as you move to the right because electrons in the outermost principal energy level experience a greater Z eff

Practice Choose the element with the higher first ionization energy (a)Al or S(b) As or Sb(c) N or Si(d) O or Cl S As N Cannot tell Page for explanations… remember trends are not reasons

Exceptions Elements would rather have one electron in each orbital than one doubled up (and not all doubled up)

Electron Affinity **Trend #4** Electron Affinity (EA) of an atom or ion is the energy change associated with the gaining of an electron by an atom in the gaseous state How much they ‘want’ an electron…

In a Nutshell… Most groups do not show a ‘trend’ But, group IA electron affinity becomes more positive as you move down the column (adding an electron becomes less exothermic) EA generally becomes more negative (adding an e- becomes more exothermic) as you move to the right across a period (or row) of the table

Metallic Character **Trend #5 Just as it sounds – How ‘good’ it is at being a metal Good conductors of heat and electricity Malleability Ductile Shiny Tend to lose electrons in a rxn

Metallic Character As you move to the right across the table, the metallic character decreases As you move down a column (or family) in the periodic table, metallic character increases

Practice Choose the more metallic element from each (a)Sn or Te(b) P or Sb(c) Ge or In(d) S or Br Sn Sb InCannot tell Page 350 – 351 for full explanations

Conceptual Connection 8.5 Try this problem on page 351 without looking at the answer Validate it using your knowledge on the trends (great test quality question)

Alkali Metals Their single valence electron (which prevents them from the noble gas configuration) is easily removed Configurations: ns 1 in the outer shell (valence shell)

Alkali Metals

Halogens These need one more electron to obtain their noble gas configuration, making them the most reactive!! Configurations: ns 2 np 5 in the outer shell (valence shell)

Halogens

Noble Gases The noble gases are the ‘happiest’ on the table… they have already obtained the octet and are stable. – Least Reactive – INERT!! Configurations: ns 2 np 6 in the outer shell (valence shell)

Noble Gases

The high ionization energies and full outer quantum levels make them very unreactive Before the 1960s, no noble gas compounds were known Since then, two of the noble gases have been shown to react with fluorine (the most reactive nonmetal on the table), under extreme conditions

Noble Gases Kr + F 2  KrF 2 Xe + F 2  XeF 2 Xe + 2 F 2  XeF 4 Xe + 3 F 2  XeF 6