Topic : Periodic Trends

Transcription

1 Topic : Periodic Trends Essential Ideas: 3.1: The arrangement of elements in the Periodic Table helps to predict their electron configurations 3.2: Elements show trends in their physical and chemical properties across periods and down groups

2 Types of Elements Metals good conductors of heat and electricity; mostly solid at room temperature; VERY diverse properties; Brittle or firm; Many have luster (shiny) Nonmetals poor conductors of heat and electricity; many are gases at room temperature Metalloids/Semiconductors possess properties of metals and nonmetals; all are solid at room temperature; semiconductors of electricity; used in electronics

3 The Periodic Table Organized from left to right by atomic number Vertical columns of the p.t. are called Groups or Families Contain elements with similar chemical properties Horizontal rows of the p.t. are called Periods As the elements fall into columns based on their properties, certain trends develop.

4 Key Words/Phrases Principal Energy Level Valence Electrons Valence Electron Energy Level Valence Electrons attracted to positive protons in the nucleus Nuclear Charge Effective Nuclear Charge Repulsion between Electrons Full Valence Electron Shell Mutual Repulsion Electron Shielding

5 Nuclear Charge Negatively charged electrons are attracted to the positively charged nucleus Related to the number of protons in the nucleus - as proton # increases in an atom, the nuclear charge also increases (becomes more positive)

6 Nuclear Charge Trend General Trend: Groups: L R Nuclear Charge Increases Periods: Top Bottom- Nuclear Charge Increases

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8 Shielding The more energy levels/interior electrons an atom has, the more shielded the outer electrons are to the nuclear charge Energy levels act like Walls to block nuclear charge from the valence electrons - makes them easier to lose (less attraction!)

9 Shielding Trend General Trend: Groups: L R Shielding is Constant (no change to interior electrons) Periods: Top Bottom- Shielding Increases with more energy levels

10 Effective Nuclear Charge Charge experienced by the valence electrons Less than nuclear charge due to shielding Effective Nuclear Charge = Nuclear Charge (#p) - Interior e- (total e- - ve- ) Energy levels act like Walls to block nuclear charge from the valence electrons - makes them easier to lose (less attraction!)

11 Effective Nuclear Charge Calculation

12 Effective Nuclear Charge General Trend: Groups: L R Increases (more protons, same # shielding e- ) Periods: Top Bottom- Constant (more protons, more shielding electrons from more energy levels)

13 Explaining ENC SHORT, CONCISE, KEY WORD-RICH ANSWERS Explain why ENC increases across a period. ENC increases across a period because the nuclear charge increases due to an increase in positive protons while shielding negative electrons remain the same, causing more attraction between the nucleus and valence electrons. Explain why ENC remains constant down a group. ENC remains constant down a group because there are more principal energy levels adding more negatively charged shielding electrons. Though there are more positive protons, the shielding electrons block the attraction from the ve- to the nucleus

14 Atomic Radius The atomic radius is ½ the distance from the nucleus of one atom to the nucleus of another, identical atom when bonded. (ex. O2) High ENC = Small Radius High Nuclear Charge with Few Shielding Electrons Strong Attraction between the nucleus and valence e- pulls the ve- in close

15 Atomic Radius General Trend: Groups: L R Decreases (Increased ENC = strong attraction) Periods: Top Bottom- Increases (More energy levels w/ ENC remaining constant)

16 Atomic Radius - Trend

17 Explaining Atomic Radius Explain why the radius of F is smaller than Lithium. F and Li have the same number of principal energy levels meaning they have the same amount of shielding. F, however, has a stronger attraction between the nucleus and the valence electrons due to its increased number of positive protons (higher nuclear charge) Explain why Ra has a larger radius than Be. Ra has 7 principal energy levels while Be has only 2. Ra has more low energy level shielding electrons blocking the attraction to the nucleus.

18 Ionic Radius Cations: Lose ve- lose an energy level Cations smaller than parent atom Anions: gain ve- increase e- repulsion in valence energy level pushing it out larger Anions larger than parent atom

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20 Ionic Radius General Trend: Groups: Decrease L R (increasing ENC = attraction) 1-14 Cations 14 smaller than : Anions Periods: Top Bottom (Increasing = more energy levels)

21 Ionic Radius - Trend

22 Explaining Ionic Radius Arrange Cl, Cl+1 and Cl-1 by increasing radii. Explain your reasoning. Cl+1, Cl, Cl-1 Cl-1 is the largest because it gained 1 electron forming a full valence energy level. The full energy level has more repulsion between the electrons, causing the radius to increase. Cl+1 has the smallest radius because it lost 1 electron, removing 1 energy level. This reduces the number of shielding electrons and increases the attraction to the nucleus. The attraction from the positive protons in the nucleus pull the ve- in close, causing a small radius.

23 Ionization Energy A measure of the attraction between the nucleus and the outer electrons M(g) M+1(g) + e The energy required to remove one mole of electrons from one mole of gaseous atoms in their ground state IE increases with an increase in ENC (strong attraction = large amount of energy to remove an electron)

24 Ionization Energy General Trend: Groups: L R Increases (more protons, same # shielding e- ) Periods: Top Bottom- Decreases (constant ENC, more energy levels = greater distance b/w p and e- = less attraction)

25 Explaining IE Explain the IE trend Across a Period. IE increases across a period because the effective nuclear charge increases. Positive protons are increasing while shielding negative electrons remain the same, causing more attraction between the nucleus and valence electrons. Explain the IE trend down a Group. IE decreases down a group because the electrons are further away from the nucleus and electrons in the lower energy levels are blocking the attraction causing electron shielding.

26 Ionization Energy - Trend

27 Exceptions to IE Red Dip (s2p1 electron configuration): the lone electron in the p orbital is easier to remove than a completely full s orbital

28 Exceptions to IE Blue Dip (s2p4 electron configuration): the 4th electron in the p orbital is easier to remove because it is being repelled by the other electron in the orbital.

29 Explaining IE Exceptions What causes B to have a lower IE than Be? B has a lower IE than Be because the 2p electrons are slightly higher in energy than the 2s electrons, so the ionization energy for B is lower than for Be What causes O to have a lower IE than N? O has a lower IE than N because the px, py, and pz only contain one electron. The extra electron in O causes a pair and repulsion making it easier to remove, hence giving O a lower IE than N.

30 Electron Affinity Opposite of Ionization Energy M(g) + e- M-1(g) The energy change when one mole of electrons are added to one mole of gaseous atoms. Generally exothermic (energy released) High Affinity = release large amount of energy Exothermic = negative number for energy Endothermic = positive number for energy

31 Electron Affinity General Trend: Groups: L R Increases (more protons, same # shielding e- ) Periods: Top Bottom- Decreases (constant ENC, more energy levels = greater distance b/w p and e- = less attraction)

32 Exceptions to Electron Affinity Electron Affinity decreases (+ energy, endothermic) with increased repulsion and decreased attraction (shielding)

33 Exceptions to Electron Affinity Explain why Beryllium (red circle) has an endothermic/very low electron affinity? Be must have an electron placed into the 2p orbital which is further from the nucleus and so experiences less attraction due to shielding from electrons in the 2s orbital. Explain why Nitrogen (blue circle) has an endothermic/very low electron affinity? The added electron must occupy a p orbital that is already singly occupied. The attraction between the electron and the atom would be reduced due to repulsion from the other p electrons.

34 Exceptions to Electron Affinity Explain why Neon (yellow circle) has an endothermic/very low electron affinity? Neon has a full valence energy level. Any additional electron must be placed in the next energy level which is farther from the nucleus. This added electron will experience increased shielding and decreased attraction to the nucleus.

35 Electronegativity The ability of an atom in a compound to attract electrons from other atoms in a covalent bond High EN = strong electron pulling power Low EN = weak electron pulling power

36 Electronegativity - Trend

37 Electronegativity Practice Predict what will occur if a very high electronegativity atom gets close to a very low electronegativity atom. An atom with high EN will have a strong attraction to the electrons of another atom while the low EN atom will not. The strong EN atom will pull the electrons from the weak atom close to it in a polar covalent bond. A large enough difference in electronegativity will cause the strong EN atom to completely remove the e- (ionize) the weak EN atom and form an ionic bond.