The Modern Periodic Table

1. The Modern Periodic Table Trends

2. Agenda Lesson: PPT, Handouts: 1.PPT Handout; Periodic Table Puzzle; Periodic Table Worksheet Text: 1. P. 30-33; Organization of the periodic table HW: 1. P. 33 # 1-4, 6; Finish all the worksheets

3. The Modern Periodic Table 1. An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column (or group). • Groups: vertical columns (#1-18) • Periodic: horizontal rows (# 1-7) 2. Periodicity – the similarities of the elements in the same group is explained by the arrangement of the electrons around the nucleus.

4. The s-block Elements: Groups 1ns1 1. Group 1: Alkali metals • soft silvery metals • most reactive of all metals, never found free in nature • reacts with water to form alkaline or basic solutions store under kerosene • whenever you mix Li, Na, K, Rb, Cs, or Fr with water it will explode and produce an alkaline solution • ns1(ending of all electron configurations for this group)

5. The s-block Elements: Group Alkaline earth metals- ns2 • less reactive than Alkali, but still react in water to produce an alkaline solution • never found free in nature • harder, denser, stronger than alkali • ns2(ending of all electron configurations for this group), because they have 2 electrons in the s sublevel, this makes them a little less reactive then the Alkali metals in group 1.

6. The d-Block Elements: Groups 3-12 • are all metals with metallic properties (malleability, luster, good conductors, etc…); are referred to as the Transition Metals • Harder and denser than alkali or alkaline • Less reactive than alkali or alkaline • For the most part their outermost electrons are in a d sublevel • Exceptions to the electron configuration are found in these groups (Ex: Ni, Pd, Pt)

7. The p – Block Elements: Groups 13 – 18 -np • Contain metals and nonmetals • Metalloids, along zigzag line, have characteristics of both metals and nonmetals (many are good conductors but are brittle). The metalloids are boron, silicon, germanium, arsenic, antimony, and tellurium.

8. Group 17 - Halogens – most reactive nonmetals-np5 • 7 electrons in outermost (s and p) energy levels (that is why so reactive – only need one electron to have 8) • called the salt formers (they react vigorously with metals to form salts). A salt is a metal and a nonmetal bonded together. • most are gases

9. Group 18 - Noble gases –unreactive-np6 • 8 electrons in outermost s and p energy levels • all are gases • The s and p blocks are called the main group or representative elements!

10. The f-Block Elements: Inner Transition Metals • final electrons fill an f sublevel • Lanthanides – shiny reactive metals; Ce-Lu (fill the 4f sublevel) • Actinides – unstable and radioactive; Th-Lr (fill the 5f sublevel)

11. Hydrogen and Helium - Oddballs • Hydrogen is NOT an Alkali metal, it is a very reactive gas. It is placed with the Alkali metals because 1s1 is its electron configuration. • Helium is a Noble gas, it is unreactive, but it does not have 8 electrons in outermost energy level, because it only has 2 total electrons!

12. # of Valence Electrons Group # Ending Configuration 1 1 ns1 very reactive 2 2 ns2 3 13 np1 4 14 np2 5 15 np3 6 16 np4 7 17 np5 very reactive 8 18 np6 very unreactive

13. Agenda Trends- Definitions • Lesson: PPT • Handouts: 1. Propeties of Atoms, • Text: 1. P. 36-41-Trends parameters and definitions • HW: 1. Finish all the worksheets

14. Graphing Assignment

15. ATOMIC RADIUS • Atomic radius is the distance from the centre of the nucleus of an atom to the outermost electron. • The greater the number of energy levels the greater is the distance of the outermost electron to the center of its atom’s nucleus. • Ionic radius is the distance from the centre of the nucleus of an ion to the outermost electron. Cationswill have a smaller ionic radius than the neutral atom. Anions will have a larger ionic radius than the neutral atom.

16. FORCE OF ATTRACTION • The force of attraction between negatively charged electrons and the positively charged nucleus is the electrostatic attraction of opposite charges. • The force of attraction existing between the outermost electron and the middle of the nucleus is dependent on two factors: 1. The size of the positive charge - determined by the number of protons in the nucleus. 2. The distance between the outermost electron and the nucleus. • A balance exists between the attraction of the electrons to the nucleus and the repulsion of the electrons between themselves

17. The size of an atomic radius cannot be measured exactly because it does not have a sharply defined boundary. However the atomic radius can be thought of as ½ the distance between the nuclei of identical atoms joined in a molecule. Trend in Atomic Radii Group trend - atomic radii decrease as you move up a group. Period trend – atomic radii decrease as you move across a period. Atomic Radii Trend

18. Example: Which is larger? P atom or Cl atom Example: Which would be larger? K+1 or K Example: Which would be larger? K+1 or Ge+4

19. Graphing Assignment

20. Ionization Energy is the energy in kilojoules per mole (kJ/mol) needed to remove the outermost electron from a gaseous atom to form a positive ion (cation) Na + Energy Na+ + e- neutral sodium has 11 protons and 11 electrons removal of 1 electron leaves 10 electrons and 11 protons and a net imbalance of charge of +1 NOTE: Metals react to lose electrons The stronger an electron is held the greater the IE needed to ionize (pull away) that electron Ionization Energy (IE)

21. Successive Ionization Energy • After the outermost electron (First IE) is removed the successive ionization energies (Second and Third IE and so on) increase as it becomes more difficult to remove the next electrons since the pull of the nucleus becomes stronger and electrons are more tightly held.

22. Group trend – ionization energy increases as you move up a group (or decreases as you move down a group). Period trend – ionization energy increases as you move across the period. Trend: Ionization Energy (IE) • Which atom has the higher first ionization energy? • Hf or Pt • Pt • (B) Cl or Ar • Ar Highest IE Trend

23. ELECTRON AFFINITY [EA] • Electron affinity is the energy released in kilojoules per mole (kJ/mol) when an electron is captured by an atom to form a negative ion (anion) Cl + Electron Cl- + Energy • neutral chlorine has 17 protons and 17 electrons • addition of 1 electron gives 18 electrons and 17 protons and a net imbalance of charge of -1 • NOTE: Nonmetals react to gain electrons

24. Period trend – electron affinity increases as you move across a period because atoms become smaller and the nuclear charge increases. This means there is a greater pull from the nucleus. Group trend – electron affinity increases as you move up a group (or decreases as you move down a group) because the size of the atom increases. Example: Which element has the greater electron affinity? Pbor Sn Sn Trend: Electron Affinity (EA) Highest EA Trend

25. Electron affinity vs. Ionization energy Electron affinity and Ionization energy follow the same trend in the periodic table. • The stronger the attraction an atom has for electrons the harder it will be to remove electrons from that atom and the higher the IE energy will be. •  The greater the attraction for electrons the greater the energy released when an atom gains an electron.

26. ELECTRONEGATIVITY [EN] • Electronegativity is a measure of the tendency of an atom to gain electrons when it is chemically combined (bonded) to another element. • The stronger the ‘pull’ or attraction of electrons to an atoms nucleus, the greater its tendency to gain electrons In general, metals have low EN and nonmetals have high EN. The actual amount of EN an atom has is indicated by a number of the Pauling Electronegativity Scale that goes from 0 to 4. Dr. Linus Pauling set up this scale and gave the element having the greatest EN an arbitrary number of 4, and he assigned numbers to the others relative to this element.

27. Period trend - EN increases as you go across a period (excluding the noble gases) because size decreases. Group trend - EN increases as you go up a group because there is less pull from the nucleus as the electrons get further away. Example Which would have the greater EN? Ca or Se Se Electronegativity enables us to predict what type of bond will be formed when two elements combine. Trend: Electronegativity (EN) Highest Electronegativity Trend

28. Electronegativity Chart

29. Electronegativity Chart

30. Electronegativity Chart

31. Agenda Expanations of Trends and Summary Lesson: PPT Handouts: 1. Propeties of Atoms, 2. Trends in the Periodic Table Summary Sheet Text: 1. P. 36-41-Trends parameters and definitions HW: 1. P. 41 # 1-7; 2. Finish all the worksheets

32. Reactivity -how easily a substance reacts with another Nonmetals gain electrons ( Electron Affinity) Metals lose electrons ( Ionization Energy) Reactivity of Nonmetals Highest Metal Reactivity Trend Nonmetal Reactivity Trend Highest

33. TRENDS IN THE PERIODIC TABLE - SUMMARY SHEET

39. Effective Nuclear Charge

40. Effective Nuclear Charge

41. Effective Nuclear Charge ENC = Number of Protons – Number of Inner Electrons

42. EFFECTIVE NUCLEAR CHARGE AND SHIELDING • The force of attraction between positively charged protons in the nucleus and negatively charged electrons is the force that holds atoms together. • The inner electrons (not in the outermost energy level) in inner energy levels, partially block or shield the attraction of the protons from the outer electrons in the outermost energy level (VALENCE ELECTRONS). • The canceling of the positive nuclear charge is called SHIELDING EFFECT.

43. EFFECTIVE NUCLEAR CHARGE (ENC) • EFFECTIVE NUCLEAR CHARGE (ENC) is a number assigned to elements to describe the amount of shielding felt by the valence electrons. ENC = Number of protons - Number of inner electrons • The greater the ENC the less the valence electrons are shielded and the stronger the pull on the valence electrons. • Greater ENC will mean a smaller atomic radius. • Shielding will help explain some of the trends in the periodic table

44. Agenda Successive Ionization Energies Lesson: PPT- Take up of all the problems Handouts: 1. Properties of Atoms, 2. Trends in the Periodic Table Summary Sheet Text: 1. P. 36-41-Trends parameters and definitions HW: 1. Finish all the worksheets 2. P 47 1-18; P. 48 #1-19,31; P. 48 # 47,55-57,65-69.

45. SUCCESSIVE IONIZATION ENERGIES • The first ionization energy is the energy required to remove the outermost electron (First IE). It is relatively low because of the repulsion exerted by the other electrons   • Each successive ionization energy (Second and Third IE and so on) will increase. • It becomes more difficult to remove successive electrons since the pull of the nucleus becomes stronger (greater number of protons relative to the electrons) and the electrons are more tightly held Ionic radius becomes smaller • There will be a noticeable jump in the increase of IE once the noble gas configuration has been reached This is because outer energy level has been removed ( radius is smaller)

46. Successive Ionization Energies Example 1: Consider the following Ionization Energies for an element X: How many valance electrons does this element have?

47. ANS: The element has 2 valence electrons. Removing the third electron from X 2+ involves a much greater energy. The third electron is closer (one energy level closer) to the attracting nucleus since the noble gas configuration has been reached.

48. Example 2: Where would the large increase in I.E. occur for Se? Explain your answer. The large increase would occur going form 6th to 7th IE. There is a noticeable jump in increase of I.E. since the noble gas configuration has been reached. (Se 6+) 1s2 2s2 2p6 3s2 3p6