Ch. 6 Chemical Periodicity

effective nuclear charge _______________________(from the protons in the nucleus) Attraction for a valence electron is partially shielded by the inner shell electrons. The simple way to estimate effective nuclear charge is to start with the atomic number (total protons) and subtract the number of inner shell electrons. Hydrogen has an effective nuclear charge of 1 ( ) Helium is 2 ( ) Lithium is 1 ( ) Sodium is 1 ( ) All transition metals are 2. All halogens are 7. Chlorine ( ), Bromine ( )

Periodic trends Many properties of elements vary with their position on the Periodic table. Various properties that can help predict chemical properties and chemical bonding are Atomic Radii Ionic Radii Ionization Energy Electron Affinity Electronegativity

Atomic radius Atomic radii decreases from left to right and increases from top to bottom. - as you move across a period, the atomic radius goes down b/c more e- = higher nuclear charge, also e- on the same n orbital (energy level) are more attracted to the (+) nucleus - as you go down a family, the radius increases b/c of higher values of n.  Which is bigger? Ge or Br F or I P or Sn

Ionic Radius – as compared to the neutral element - Positive _______ are always smaller than neutral atom (removal of e-) Na 1s22s22p63s1 Na+ 1s22s22p6 often remaining e- are in a lower n   - _________ get bigger than neutral atom. Adding e- creates more e- repulsion. Which is bigger? Cl or Cl- F- 10e-, 9p+ or Ne 10e-, 10p+ K+ 18e-, 19p+ or Cl- 18e-, 17p+

Ionization Energy I.E. _________________, the energy needed to remove the outermost electron. The greater the ionization energy, the stronger the attraction and the more energy needed to remove that electron. IE is measured as the amount of energy required to remove an e- from an isolated gaseous atom. It measures how tightly atoms hold on to their electrons Na 11e-  Na+ 10e- -Atoms can have 1st I.E., 2nd I.E., 3rd I.E., etc. The 2nd I.E. is the amount of energy needed to remove the 2nd e- …

In general, 1st I.E. increase as you move towards the right and decrease as you move down a group in the periodic table. - Across a period, I.E. goes up: Metals – like to lose e-. Low I.E. Nonmetals – like to gain e-, close to a full valence, High I.E. - Down a family, I.E. goes down: much easier to take away an e- b/c lots of e-, more repulsion, very low attraction, shielding.

- Successive I.E.’s for the same atom go up. Al  Al+  Al2+  Al3+  Al4+ Why? For Al: IE1 < IE2 < IE3 < < (much greater) IE4 Why?

Which has a higher I.E.? Si or S S or F O, Cl, or H Observing the 3 I.E. for element X, what family of elements does X belong to? IE1 = 300kJ IE2 = 620kJ IE3 = 8000kJ   _____ b/c the third e- comes from an inside energy level.

Electron Affinity EA ___________________is the amount of energy absorbed when an atom gains an e- in the gaseous form to make an anion with a -1 charge. In general, EA become more negative from left to right and from bottom to top. Across a period, EA goes down; needs less energy to happen - Down a family, EA goes up, less tendency to gain e- b/c of shielding and e- repulsion.

Electronegativity χ (give me those electrons) ____________________ is an atom’s ability to attract an e- in a chemical bond (F has the highest χ = 4) A combination of IE & EA. χ usually increased from left to right and decreases from top to bottom.   χ across a period increases b/c nonmetals want to gain e- while metals would rather give up e- to have a full valence χ down a family decreases

χ can be used to make predictions about bonding. If two elements have very different χ, they will tend to make ionic bonds. The less χ element (metal) will give up its electron to the more χ element. Two nonmetals with similar χ will form covalent bonds. While sharing e- the more electronegative element will have a greater share of the e-

Summary of trends Elements with ______ I.E. and _______ χ lose e- easily to form cations Elements with ______ I.E. and ______ negative EA and high χ gain e- to form anions Cations are always ______ than the original atom Anions are always ______ than the original atom

*For all trends except size, the closer you are to F, the greater the trend. ________________ - the easier an atom gives e- the more metallic it is (metals) __________________ – all species have the same # of electrons (Na, Mg+, Al2+)

Coulomb’s Law ____________ : for electrostatic attraction/repulsion. Applies to charged particles, magnets, gravitation, ionization energy, lattice energy F = k since we are not calculating, just comparing, let’s simplify to use energy instead of force E = k E = energy of attraction or repulsion between charged particles Q1 = charge of first particle Q2 = charge of second particle r = distance between charged particles

if opposite charges - The greater (bigger) the charges the greater the ____________ if like charges - The greater (bigger) the charges the greater the ____________ Also the closer the particles are the ____________ the attraction or repulsion will be.

When Applying Coulomb’s Law to Atoms and Ions E = ionization energy Q1 = charge of an electron, -1. Q2 = effective nuclear charge of protons in nucleus r = distance between charged particles which can be approximated by the Period – two atoms in the same period have _____________________ the same atomic radius. Na (Period 3) is smaller than K (Period 4) but similar in size to Al (also Period 3.)

Using only a Periodic Table 1. Which is larger, Na or K? Explain why. ________________________. The size of an atom is determined by the energy level of its’ valence electron(s). K has 4 occupied energy levels while Na has 3. Therefore _______________________.

2. Which is larger, Na or Al. Explain why. ____________ 2. Which is larger, Na or Al? Explain why. ____________. Both Na and Al have 3 occupied energy levels, so they should be approximately the same size. Adding protons and electrons does not add to the size, since they fill vacant space in either the nucleus or the valence shell. But the attraction for the valence electron in Al (effective nuclear charge ____________) is ____________than that of Na (effective nuclear charge ____________ ) So the force of attraction is ____________ in Al and its’ valence electrons are pulled a little closer to the nucleus, giving Al a ____________ atomic radius than Na.

3. Which is larger, Na or Na+. Explain why. ____________ 3. Which is larger, Na or Na+? Explain why. ____________. Na has 11 electrons, so it’s valence electron shell is 3. Na+ has lost an electron and is _____________________ with Ne which has only 2 occupied levels which makes it ____________ than Na.

4. Which is larger, Br or Br-. Explain why 4. Which is larger, Br or Br-? Explain why. ____________ Both Br and Br- have the same values for everything in ____________________, so they are approximately the __________size. But even though the added electron in Br- goes into a vacant space in the valence shell, electrons have like charges and repel one another. This increased ____________ between electrons in the valence shell pushes them slightly apart and makes the _______________________.

5. Which has a higher first ionization energy, Na or K. Explain why 5. Which has a higher first ionization energy, Na or K? Explain why. _____________________________________________. Both have the same electron charge and effective nuclear charge. But the valence electrons for K are in the 4th energy level, while those of Na are in the 3rd. So the electrons of Na are held by a ___________________ ___________________________________. Which is why metallic character (_____________________ ____________) generally increases as you move down a family. (Or if you prefer why first ionization energy decreases as you move down a family – its’ the same thing.)

6. Which has a higher first ionization energy, Na or Al. Explain why 6. Which has a higher first ionization energy, Na or Al? Explain why. _______________________________________________ Both have the same electron charge, and the valence electrons of both are in the 3rd energy level. But Na has an effective nuclear charge of 1 (11 – 10 = 1) while Al has an effective nuclear charge of 3 (13 – 10 = 3.) So the valence electrons of Al are held ____________ and it requires _____________________ to remove them. Which is why ____________ ____________ increases as you more to the left across a Period and decreases as you move to the right. (Or why ____________ ____________ increases as you move to the right across a Period.)

Fossil fuels and pollution _____________________ are primarily hydrocarbons yet they are usually contaminated with sulfur. During combustion, sulfur dioxide is formed which is a harmful air pollutant. SO2 is oxidized to SO3 in air, which is the acid anhydride (water has been eliminated from the compound) of sulfuric acid. Sulfuric acid is then the major contributor to acid rain.

Nitrogen compounds are also impurities found in fossil fuels Nitrogen compounds are also impurities found in fossil fuels. They undergo combustion to form nitrogen monoxide. Airplanes, cars, furnaces, etc exhaust NO into the atmosphere. NO is oxidized into NO2 which is the major component of the reddish-brown haze of photochemical smog. NO2 also reacts with water to form nitric acid.

When the elements Np and Pu were discovered by McMillan and Seaborg, they were placed on the periodic table just below La and Hf. Yet, after studying the chemicals for a few years, Seaborg decided they should be placed in a new row below the lanthanides. What justifications could he have used to move these elements. What do the catalytic converters on our cars exhaust systems actually do? How do they decrease air pollution?