Periodic Trends This document discusses several periodic trends including ionization energy, electronegativity, atomic radius, and ionic radius. It notes that ionization energy generally decreases down a group and increases left to right in a period as the nuclear charge increases. Electronegativity also typically increases left to right and decreases down a group. Atomic radius decreases left to right in a period as the nuclear charge increases but increases down a group as the principal energy level increases. Ionic radius trends are similar, with positive ion size decreasing left to right and both positive and negative ion size increasing down a group.

1. Periodic Trends Chapter 6 Section 3 pp.163-169

2. What is a trend? 1. The general direction in which something tends to move. 2. A general tendency or inclination. See Synonyms at tendency.

3. Things to remember… An element’s # of protons does not change Isotopes- atoms of an element with different masses. Different number of neutrons Nuclear charge- over all charge obtained from the attraction of all the protons and all e- in an atom Positively charged nucleus pulls electrons towards it

5. Ions An atom or bonded group of atoms that has a positive or negative charge Occurs when an atom gains or loses an electron Atoms that gain e- have a negative overall charge (anion) Atoms that lose e- have a positive overall charge (cation) Electrostatic repulsion Remember like charges repel each other…electrons in orbitals repel/push each other away

7. Ionization Energy Ionization energy (IE) is the energy required to remove an electron from a gaseous atom (kJ/mol) Energy is needed to overcome the attraction between the positive protons and negative electrons IE indicates how strongly a nucleus can hold onto its valence e- High IE= strong hold on e-; less likely to make positive ions Low IE= atom can lose valence e- easily; likely to make positive ions

8. Octet Rule What is the electron configuration for sodium? What is the new configuration when a sodium atom becomes a +1 sodium atom? What noble gas does this configuration look like? OCTET RULE: Atoms tend to gain, lose, or share electrons in an order to acquire a full set of 8 valence electrons First period elements are en exception to the rule (2Ve-) This rule helps to determine what kind of ions will form

9. More about ionization energy… 1st Ionization Energy (1st e-) After removing the 1st e- it is still possible to remove additional electrons Second ionization energy: The IE needed to remove a second electron from a +1 ion For sodium IE will be high since it does not want to lose anymore e-, reached 8 valence e- Third ionization energy: The IE needed to remove a third electron from a +2 ion 1e- 8e- 11p+ 12n 2e- Sodium

10. Ionization Energy What type of IE does group 1A have? High or low? Low IE Likely to form + cations What type of IE does group 8A have? High IE Unlikely to form cations From left to right, the IE to remove successive electrons always increases Does not happen smoothly…requires large jump of energy

11. IE continued…. Trends in periods: First IE generally increases as you move left-to-right The nuclear charge increases (greater attraction) with each successive element… What does this have to do valence electrons? Increase the nuclear charge=stronger hold on valence electrons

12. IE continued… Trends within groups… First IE generally decrease as you move down a group What happens to the size of an atom as you move down a group? Increases This cause decrease in IE Valence electrons are farther from the nucleus Easier to remove

13. Electronegativity The ability of atoms of an element to attract electrons in a chemical bond Fluorine is the most electronegative element Value: 3.98 Cesium and Francium are the least electronegative Values: 0.79 and 0.7 Atom with greater electronegativity attracts electrons in a chemical bond the strongest Leave out noble gases (they really don’t react)

15. Electronegativity (EN) Trends in periods and groups: EN decreases as you move down a group Increases as you move left to right on table Where are the lowest electronegatives found? Lower left side of the periodic table Where are the highest electronegatives found? Upper right side of the periodic table

17. Electron Affinity The energy change that occurs when an electron is acquired by a neutral atom Measure in kJ/mol High Eaffinity = more negative the nrg, easier to gain e- Low Eaffinity = more positive the nrg, easier to lose e- Noble gases have “zero” Eaffinity

18. Electron Affinity Many atoms release energy when they gain an electron A + e-  A- + energy A + e- + energy  A- Some atoms need energy to be “forced” to gain an electron Produces unstable ions that lose electron quickly

20. Food for thought…. Do atoms on the right side of the periodic table tend to gain or lose electrons? Gain electrons What type of ions are they likely to form? Negative ions Do atoms on the left side of the periodic table tend to gain or lose electrons? lose e- What type of ions do they form? Positive ions

21. Atomic Radius Electron cloud Spherical surface in which there is a 90% probability of finding an e- Not physical Atomic size is defined by how closely an atom lies to its neighboring atom Each atom has different properties so there are going to be different sizes… Different sizes in different blocks…

23. Atomic Radius of a Metal Metals Atomic radius half the distance between adjacent nuclei in a crystal of that element

24. Atomic Radius of a Nonmetal Usually occur as molecules Atomic radius: Half the distance between nuclei of identical atoms that are chemically bonded together Diatomic molecules Hydrogen Bromine Iodine Oxygen

25. Atomic Radius: Trends in Periods DECREASE in atomic radii left-to-right Increasing positive charge in the nucleus Principle energy level (n) remains the same throughout the period Each successive element the atomic # Increase Add proton and electron Electron gets added to same Prin. E level (n) No additional electrons come across valence electrons and nucleus Val. E- are not shielded from increased nuclear charge Therefore the increased nuclear charge brings the outermost electrons closer to the nucleus …which means… Smaller atomic radii going   

26. Atomic Radius: Trends within Groups INCREASE as you move down a group Nuclear charge increases What happens to your principle energy level as you move down a group? Therefore…Electrons added to higher principle energy levels Although nuclear charge increased, other factors in play to overpower increased nuclear charge... Outermost orbital increases in size Electrons are farther from the nucleus b/c of bigger orbital More resistant to higher nuclear charge b/c of increased distance Principle energy level increased Puts more orbital with electrons between the nucleus and the outermost electrons These electrons in between shield the outermost electrons from the pull of the nucleus

28. Solve… Which has the largest radius? Magnesium (Mg) Silicon (Si) Sulfur (S) Sodium (Na) Which has the smallest radius?

29. And the answer is… Largest: Na (sodium) Smallest: S (sulfur)

30. Ionic Radius Losing/gaining electrons affects size of an atom Gain an electron atom becomes negative atom becomes larger Electrostatic repulsion between atom’s outer electrons increase Forces outer electrons to move further apart Makes radius bigger Lose an electron atom becomes positive atom becomes smaller Electron lost is valence electron Can lead to completely empty orbital=smaller radius Electrostatic repulsion between atoms decrease so they can be pulled closer to nucleus

31. Ionic Radius: Trends in Period What type of ions do you think will be formed on the left side of the table? Smaller positive ions What type of ions do you think will be formed on the right side of the table? Larger negative ions General Rule: Left-to-right across a period, the size of positive ions gradually decreases Around 5A and 6A, the size of much larger negative ions gradually decreases

33. Ionic Radius: Trends in Group What happens to the principle energy level as we move down a group? Increases Ion’s outer electrons are in a higher principle level This means an increase in ionic size Ionic radii increases as we move down a group for both positive and negative ions