Periodic trends

1. Periodic trends Gradual changes -> or Large jumps 

2. Trends • The Periodic Table is arranged so that predictions can be made about the elements’ chemical behavior.

3. Periods and Group • Period – horizontal row on P.T. • Each period represents an energy level (think back to models of the atom) • Atoms in period 1 have 1 energy level, atoms in period 5 have 5 energy levels • Group – vertical column P.T. • Each group represents a certain number of valence electrons • Also known as families

4. Valence Electrons • Electrons exist within orbitals. • The electrons in the outermost energy level are called valence electrons! • Valence electrons determine how an atom behaves (they perform all bonding)

5. Valence Electrons Ctd • Each group number represents the number of valence electrons elements in that group have • Ex: group 1 has 1 valence electron • No atom can have more than 8 v.e. • For group numbers 13 - 18 subtract ten to figure the number of valence electrons • Ex: Group 18 = 8 valence electrons • Groups 3-12, we assume to have 2 v.e. (not always the case)

6. Valence Electrons

7. 8 is great (sometimes 2) • The noble gases (group 18) are stable atoms, meaning they do not react! (this is good) • The noble gases are stable because they have 8 valence electrons (or 2 as in helium) • All other atoms will gain or lose electrons to become like noble gases (remember cations and anions)

8. Metals Form Cations • Groups 1-13 are metals (except H) • Metals lose electrons to become like noble gases (less than 4 v.e., lose e-) • Therefore, they form cations (+ charge)!

9. Non-Metals Form Anions • Groups 14 – 18 are non-metals • They gain electrons to be like noble gases! (more than 4 v.e., gain e-) • Therefore, they form anions (- charge)

10. Practice • For each of the following elements, determine if they are a metal or non-metal and the charge they would form. Na O F Al Ba I S Cs P

11. Periodic Table Families • Group 1: Alkali Metals • Group 2: Alkaline Earth Metals • Groups 3-12: Transition Metals • Group 17: Halogens • Group 18: Noble Gases

12. Shielding • As atom gets bigger, more E levels are added • Therefore, valence electrons get farther from pull of nucleus • Shielding occurs when the inner electrons block the positive pull from the nucleus from reaching the valence electrons • Only impacts trends down a group!

13. Electron Shielding

14. Things to Consider • P+ and e- within an atom are attracted to each other, but P+ pull e- because e- have less mass • As you go down a group, each atom adds an energy level! • As you go across a period each element adds a p+ and e-! (we consider all elements on P.T. neutral)

15. Atomic Radius – size of atom • What do you think? • Does atomic radius increase/decrease across a period? • Why? • Does atomic radius increase/decrease down a group? • Why?

16. Atomic Radii Visual

17. Atomic radius (size) • A.R. increases down a group (atom gets bigger) • Add energy levels (electrons farther from nucleus) • Shielding occurs • Fr is bigger than H

18. Atomic Radius (cont) • Decreases as you move across a period (atom gets smaller) • Adding more protons, so outermost electrons are pulled closer • Ex: K is bigger than Br

19. Determine which atom has a larger A.R. • Be vs. Mg • Co vs. Ir • Re vs. At • Mn vs. Zn • He vs. Ar • B vs. Ga Mg, more energy levels & more shielding Ir, more energy levels & more shielding Re, less protons Mn, less protons Ar, more energy levels & more shielding Ga, more energy levels & more shielding

20. Ionization Energy (I.E.) • Energy required to remove an outermost electron (valence e-) • What do you think? • Does I.E. increase/decrease across a period? • Why? • Does I.E. increase/decrease down a group? • Why?

21. Ionization Visual

22. Ionization Energy (I.E.) • I.E. decreases down a group (easier to take electron) • more energy levels, therefore valence e- far from the attractive force of the nucleus • Electron shielding occurs down a group • Ex: I.E. lower for Fr (francium) than H (hydrogen) • Stealing a basketball/Attacking Gazelle

23. Ionization Energy (cont) • I.E. increases across a period (harder to take electron) • More protons being added, pulling outer electrons in more! • Atoms are getting closer to being like a noble gas (do not want to give up e-!) • Ex: Kr (krypton) requires lower IE than K (potassium)

24. Determine which atom has a larger I.E. • Ca vs. Ba • Mg vs. P • O vs. Se • Mo vs. Pd • H vs. Li • K vs. V Ca, less energy levels, smaller atom P, closer to being Noble gas, more protons O, less energy levels, O smaller atom Pd, more protons to hold electrons tighter H, less energy levels, smaller atom V, more protons to hold electrons tighter

25. Electronegativity • Ability of an atom to attract an electron • Low electronegativity = hard to attract an e- • High electronegativity = easy to attract an e- • What do you think? • Does E- increase/decrease across a period? • Why? • Does E- increase/decrease down a group? • Why?

26. Electronegativity Visual

27. Electronegativity • E.N. decreases down a group (harder for atom to attract electrons) • E- shielding occurs • More energy levels so distance between nucleus and electrons increases • Ex: H has higher electronegativity than Fr (H has easier time attracting e- than Fr)

28. Electronegativity (cont.) • E.N. increases as you move across a period • Nuclear charge increases (more p+) across a period causing e- to be attracted much more strongly • Metals do not want to gain e-, while nonmetals add e- to become like noble gas! • Ex: K has lower electronegativity than Br (K has a harder time attracting e-)

29. Electronegativity Chart

30. Determine which atom would have higher electronegativity • Be vs Ca • N vs O • F vs Cl • C vs. Pb • F vs. Li • Sc vs. Co Be, less energy levels, smaller atom, less shielding O, closer to being Noble gas, more protons F, less energy levels, smaller atom, less shielding C, less energy levels, smaller atom, less shielding F, closer to being Noble gas, more protons Co, more protons to attract e- more

31. Reactivity • How easily an element will gain or lose electrons • Focuses on metal & non-metals in P.T.

32. Reactivity • What metals do you think are most reactive? Least reactive • Why? • What non-metals do you think are most reactive? Least reactive? • Why

33. Reactivity Visual

34. Reactivity – Down Group • Metals – increases (Li less reactive than Fr) • Metals want to lose electrons to gain stability • Easy to lose electrons as you add energy levels as electrons are further from pull of nucleus • Non-Metals – decreases (F more reactive than I) • Non-metals want to gain electrons to gain stability • Hard to gain electrons as you add energy levels as electrons are farther from pull of nucleus

35. Reactivity – Across Period • Metals – decreases (K more reactive than Ca) • Metals want to lose electrons to have a full octet • as you move across period, metals are adding electrons & protons • Non-metals – increases (N less reactive than F) • Non-metals want to gain electrons to have a full octet • As you move across period, non-metals are adding electrons

36. Determine which atom is more reactive • Be vs. Mg • Na vs. Al • O vs. S • P vs. Cl • Li vs. Cs • N vs. F Mg, valence e- further from nucleus, bigger Na, has to lose 1 valence e- O, valence e- closer to nucleus, less energy levels Cl, has to gain 1 valence e-, more p+ Cs, valence e- further from nucleus, bigger F, has to gain 1 valence e-, more p+

37. Tips for Trends • Focus on if element wants to gain or lose e- to become stable • Re-state trend in your own words • Know what shielding does! • Do not try to memorize!