Periodic Properties of the Elements

Transcription

1 Chapter 7 Periodic Properties of the Elements DEVELOPMENT OF THE PERIODIC TABLE Elements in the same group generally have similar chemical properties. Properties are not identical, however. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

2 Alkali Metals Periodic Table Noble Gases Alkaline Earths Halogens Main Group Transition Metals Main Group Lanthanides and Actinides Prentice-Hall 2002 General Chemistry: Chapter 10 2

3 Electronic Configuration 1 3

4 Effective Nuclear Charge In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. The nuclear charge that an electron experiences depends on two factors. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

5 Effective Nuclear Charge, Z Eff Net positive charge experienced by an electron Z eff = Z - S Z = number of protons in the nucleus S = average number of electrons responsible for screening In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. The nuclear charge that an electron experiences depends on both factors above. e.g. Na = 1s 2 2s 2 2p 6 3s 1 [Ne] 3s 1 5

6 Periodic Trends Trends in the following experimentally determined properties: Atomic Radii Ionic Radii Ionization Energies Electron Affinities 6

7 Sizes of Atoms and Ions The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

8 Sizes of Atoms Bonding atomic radius tends to: decrease from left to right across a row due to increasing Z eff increase from the top to the bottom of a column due to increasing value of n Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

9 Atomic Radius Trends: Atomic Radius Atomic radius Down the Group: size as n Across the period (L R): As more e - s added to same shell inefficient shielding Z eff & size 9

10 Screening & Penetration Screening - the inner electrons block the other outer electrons from the nucleus effect Penetration - is to get close to the nucleus attraction nucleus electron repulsion attraction nucleus Core electrons repulsion Valence electron 10

11 Screening & Penetration Atomic Radius decreases across the Period from L R: As we add electrons in one by one (new element) The electrons go into same shell (orbital) and do not screen the nucleus any further than before. As nuclear charge goes up, electrons all held or attracted more closely to the nucleus. Atomic Radius increases down the Group: When adding a full electronic orbital this gives more screening therefore electrons are held more loosely. 11

12 Sizes of Ions Ions increase in size as you go down a column Due to increasing value of n. In an isoelectronic series, ions have the same number of electrons. Ionic size decreases with an increasing nuclear charge. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

13 Ionic Radius The Size Changes When Atoms form Ions: Cations: Anions: p + > e - Z eff & size p + < e - e - - e - repulsion & size Trends within ions For isoelectronic cations: the greater the positive charge, the smaller the ion For isoelectronic anions: the greater the negative charge, the larger the ion 13

14 Cationic Radius Cations are smaller than their parent atoms The outermost electron is removed (Radial probability density reduced) Repulsion forces between electrons are reduced. 14

15 Anionic Radius Anions are larger than their parent atoms Electrons are added (Radial probability density is increased) Electron repulsion forces between electrons is increased Down the group - increasing value of n 15

16 Isoelectronic Ions Isoelectronic the same number of electrons Isoelectronic Cations: Example K + > Ca 2+ > Sc 3+ (18 electrons) The more positive the ionic charge, the smaller the ionic radius Isoelectronic Anions: Example Cl - < S 2- < P 3- (18 electrons) The more negative the ionic charge, the larger the ionic radius 16

17 Ionisation Energy Amount of energy required to remove an electron from the ground state of a gaseous atom or ion: First ionisation energy is that energy required to remove the first electron. Second ionisation energy is that energy required to remove the second electron, etc. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

18 Ionisation Energy It requires more energy to remove each successive electron. When all valence electrons have been removed, the ionisation energy takes a quantum leap. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

19 Trends in First Ionisation Energies As one goes down a column, less energy is required to remove the first electron. Figure 6.8 For atoms in the same group, Z eff is essentially the same, but the valence electrons are farther from the nucleus. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

20 Trends in First Ionisation Energies Generally, as one goes across a row, it gets harder to remove an electron: As you go from left to right, Z eff increases. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

21 Electron Affinities Energy change accompanying the addition of an electron to a gaseous atom: Cl + e Cl E = -349 kj/mol Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

22 Trends in Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

23 1 st Electron Affinities Filled s-shell, e - goes into p-orbital Added e - e - -e - repulsion 23

24 Second Electron Affinity Second electron affinities are all positive Why? It takes energy to overcome the repulsion between an electron and the already-negative ion. e.g. O (g) + e - O - (g) O - (g) + e - O 2- (g) EA 1 = -141 kj EA 2 = +744 kj 24

25 Metal, Nonmetals and Metalloids Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

26 Metals versus Nonmetals Differences between metals and nonmetals tend to revolve around these properties. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

27 Metals versus Nonmetals Metals tend to form cations. Nonmetals tend to form anions. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

28 Metals Tend to be lustrous, malleable, ductile, and good conductors of heat and electricity. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

29 Metals Compounds formed between metals and nonmetals tend to be ionic. Metal oxides tend to be basic. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

30 Reactivity of Metals Metal reactivity varies with the type of metal. Alkali metals are very reactive because it is very easy to remove the 1s 1 electron. All alkali metals react vigorously with water to produce metal hydroxide and hydrogen gas: 2M (s) + 2H 2 O (l) -> 2MOH (aq) + H 2(g) Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

31 Reactivity of Metals Alkaline earth metals are less reactive than alkali metals. Mg reacts only with steam, Be does not react with water, but others react readily with water. Reactivity tends to increase as you move down groups. Figure 6.14 Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

32 Nonmetals Dull, brittle substances that are poor conductors of heat and electricity. Tend to gain electrons in reactions with metals to acquire noble gas configuration. Substances containing only nonmetals are molecular compounds. Most nonmetal oxides are acidic. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

33 Metalloids Have some characteristics of metals and some of nonmetals. For instance, silicon looks shiny, but is brittle and a fairly poor conductor. Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e 2010 Pearson Australia

34 Diagonal Relationship The 1 st element in a Group behaves like the 2 nd element of the Group adjacent to it. Increase in Z eff Increase in size decrease in Z eff decrease in size The relationship exists because of size and charge effect. 1. Moving from L R in the periodic table, atomic size decreases 2. Moving down the group, the atoms & ions increase in size Thus on moving diagonally, the size remains nearly the same e.g. Li + = 0.76 Å and Mg 2+ = 0.72 Å 34

35 Diagonal Relationship These pairs (Li & Mg), (Be & Al), (B & Si ) etc. exhibit similar properties e.g. Boron and Silicon are both: Semiconductors form halides that are hydrolysed in water and have acidic oxides Because Their atoms / ion pairs have comparable sizes, ionisation energies, electron affinities, etc. 35

36 Periodic Trends - Summary 36

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