1. Chapter 7 Periodic Properties of the Elements
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2. Overview of Periodic Trends
Why and how show we study trends in the periodic table?
Focus on valence electrons  understand and predict reactivity of elements
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3. Overview of Periodic Trends
Trends within the periodic table can be qualitatively understood
and explained using Coulomb’s law, the shell model, and the
concept of shielding/effective nuclear charge.
Key concept:
Effective nuclear charge  effects on the valence electron(s)
a) First ionization energy
b) Atomic and ionic radius
c) Electronegativity (in bonding chapters)
d) Typical ionic charges
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4. Overview of Periodic Trends
Be able to predict and explain the trends, using Coulomb’s Law, effective nuclear charge (and the shielding effect) and atomic structure.
Other ways to describe these periodic trends incorporate the use of PES data with ionization energy, an application of Coulomb’s Law.
(more on this later)
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5. Overview of Periodic Trends
2. Effects on the properties of metals, nonmetals and metalloids
 group trends in terms of reactions
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6. Coulomb’s Law
The force between two charged particles is proportional to the magnitude of the two charges (q1 and q2), and inversely proportional to the square of the distance, r, between them.
If the charges are the same, they repel each other.
If the charges are opposite (protons vs. electrons), they attract each other.
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7. Coulomb’s Law F = k Q1Q2 r2 In terms of potential energy, E = k Q1Q2 r
Physics: Whenever you move a particle in a direction it doesn’t want to go in a field (e.g. electrical or gravitational), you give it PE proportional to the distance you move it. (from Mr. Lapp)
We will apply Coulomb’s Law qualitatively
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8. 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 both factors.
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9. Effective Nuclear Charge a way to quantify the shielding effect
The effective nuclear charge, Zeff, is found this way:
Zeff = Z − S
where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons.
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10. Shielding Effect
The position of a valence electron and the ability to remove it from an atom are related to
the number of protons in the nucleus
the extent to which the valence electron is shielded from the positively-charged nucleus by the negatively-charged core electrons
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11. What Is the Size of an Atom?
The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.
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12. Sizes of Atoms Bonding atomic radius tends to…
…decrease from left to right across a row (due to increasing Zeff).
…increase from top to bottom of a column (due to increasing value of n).
We can also describe this trend using the shielding effect.
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13. Sizes of Ions Ionic size depends upon: The nuclear charge.
The number of electrons.
The orbitals in which electrons reside.
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14. Sizes of Ions Cations are smaller than their parent atoms.
The outermost electron is removed and repulsions between electrons are reduced.
 stronger Coulombic attraction between protons and valence electrons
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15. Sizes of Ions Anions are larger than their parent atoms.
Electrons are added and repulsions between electrons are increased.
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16. Sizes of Ions Ions increase in size as you go down a column.
This is due to increasing value of n.
and the larger shield of core electrons
 decreasing Coulombic attraction
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17. Sizes of Ions
In an isoelectronic series, ions have the same number of electrons.
Ionic size decreases with an increasing nuclear charge.
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18. Ionization Energy
ionization energy = the amount of energy required to remove an electron from the ground state of a gaseous atom or ion.
The first ionization energy is the energy required to remove the first electron.
The second ionization energy is that energy required to remove a second electron, etc.
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19. Ionization Energy
It requires more energy to remove each successive electron.
When all valence electrons have been removed, the ionization energy takes a quantum leap.
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20. Trends in First Ionization Energies
As one goes down a column, less energy is required to remove the first electron.
For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.
Increased shielding effect down the group, resulting in more exposure of the valence electrons
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21. Trends in First Ionization Energies
Generally, as one goes across a row, it gets harder to remove an electron.
As you go from left to right, Zeff increases.
Same number of core electrons = same size shield as you go from left to right, BUT increasing attraction by the nucleus for the valence electrons makes it harder to remove an electron.
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22. Trends in First Ionization Energies
However, there are two apparent discontinuities in this trend.
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23. Trends in First Ionization Energies
The first occurs between Groups IIA and IIIA.
In this case the electron is removed from a p-orbital rather than an s-orbital.
The electron removed is farther from nucleus.
There is also a small amount of repulsion by the s electrons.
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24. Trends in First Ionization Energies
The second occurs between Groups VA and VIA.
The electron removed comes from doubly occupied orbital.
Repulsion from the other electron in the orbital aids in its removal.
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25. Electron Affinity
Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom:
Cl + e−  Cl−
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26. Ionization Energy vs. Electron Affinity
Ionization energy: how much does it cost to remove an electron?
Always positive (+kJ)
Electron affinity: What is the energy change when an electron is added to an atom?
Positive or negative (+kJ or –kJ)
Together ionization energy and electron affinity
electronegativity
bonding and reactivity
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27. Properties of Metal, Nonmetals, and Metalloids
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28. Metals
They tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.
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29. Metals Compounds formed between metals and nonmetals tend to be ionic.
Metal oxides tend to be basic when dissolved in water.
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30. Reactions of Metal Oxides
Most metal oxides are basic:
Metal oxide + water  metal hydroxide
(synthesis reaction)
Na2O(s) + H2O(l)  2NaOH(aq)
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31. Reactions of Metal Oxides
Metal oxides are able to react with acids to form salts and water:
Metal oxide + acid  salt + water
MgO(s) + 2HCl(aq)  MgCl2(aq) + H2O(l)
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32. Practice Exercise 7.8
Write the balanced chemical equation for the reaction between copper (II) oxide and sulfuric acid.
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33. Metals versus Nonmetals
Differences between metals and nonmetals tend to revolve around these properties.
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34. Metals versus Nonmetals
Metals tend to form cations (are oxidized), including oxides (think rust).
Nonmetals tend to form anions (are reduced).
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35. Nonmetals
These are dull, brittle substances that are poor conductors of heat and electricity.
They tend to gain electrons in reactions with metals to acquire a noble gas configuration.
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36. Nonmetals
Substances containing only nonmetals are molecular compounds.
Most nonmetal oxides are acidic when dissolved in water.
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37. Nonmetals
Seven nonmetallic elements exist as diatomic molecules under ordinary conditions:
H2(g), N2(g), O2(g), F2(g), Cl2(g), Br2(l), I2(s)
Remember the “7” rule.
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38. Reactions - Nonmetals
When nonmetals react with metals, nonmetals tend to gain electrons:
Metal + nonmetal  salt
2Al(s) + 3Br2(l)  2AlBr3 (s)
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39. Reactions - Nonmetals
Most nonmetal oxides are acidic when dissolved in acid:
Nonmetal oxide + water  acid (synthesis)
P4O10(s) + 6H2O(l)  4H3PO4(aq)
CO2(g) + H2O(l)  H2CO3(aq)
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40. Reactions - Nonmetals
Nonmetal oxides react with bases to form salts and water:
Nonmetal oxide + base  salt + water
CO2(g)+ 2NaOH(aq)  Na2CO3(aq) + H2O(l)
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41. Sample Exercise 7.9 (p. 256)
Write the balanced chemical equations for the reactions of solid selenium dioxide with
a) water
b) aqueous sodium hydroxide
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42. Practice Exercise 7.9
Write the balanced chemical equation for the reaction of solid tetraphosphorus hexoxide with water.
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43. Metalloids
These have some characteristics of metals and some of nonmetals.
For instance, silicon looks shiny, but is brittle and fairly poor conductor.
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44. Group Trends
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45. Alkali Metals Alkali metals are soft, metallic solids.
The name comes from the Arabic word for ashes.
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46. Alkali Metals
They are found only in compounds in nature, not in their elemental forms.
They have low densities and melting points.
They also have low ionization energies.
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47. Alkali Metals Reactions
Alkali metals react with hydrogen to form hydrides.
In hydrides, the hydrogen is present as H–, called the hydride ion.
2M(s) + H2(g)  2MH(s)
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48. Alkali Metals
Alkali metals react with water to form MOH and hydrogen gas:
e.g. Na(s) + H2O(l)  NaOH(aq) + H2(g)
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49. Alkali Metals Reactions
Alkali metals produce different oxides when reacting with O2:
4Li(s) + O2(g)  2Li2O(s) (oxide)
2Na(s) + O2 (g)  Na2O2(s) (peroxide)
K(s) + O2 (g)  KO2(s) (superoxide)
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50. Alkali Metals
Alkali metals (except Li) react with oxygen to form peroxides.
K, Rb, and Cs also form superoxides:
K + O2  KO2
They produce bright colors when placed in a flame.
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51. Sample Exercise 7.10 (p. 279)
Write balanced chemical equations that predict the reactions of cesium metal with
a) Cl2(g)
b) H2O(l)
c) H2(g)
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52. Practice Exercise 7.10
Write a balanced chemical equation that predicts the products of the reaction between potassium metal and elemental sulfur (S8).
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53. Alkaline Earth Metals
Alkaline earth metals have higher densities and melting points than alkali metals.
Their ionization energies are low, but not as low as those of alkali metals.
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54. Alkaline Earth Metals Reactions
Their chemistry is dominated by the loss of two s electrons:
M  M2+ + 2e–
Mg(s) + Cl2(g)  MgCl2(s)
2Mg(s) + O2(g)  2MgO(s)
(remember last year’s Reactions labs – white flame?
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55. Alkaline Earth Metals Reactivity increases down the group.
Beryllium does not react with water and magnesium reacts only with steam, but the others react readily with water.
Thursday’s lab
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56. Hydrogen
Hydrogen is a unique element, most often occurs as a colorless diatomic gas, H2.
Reactions between hydrogen and nonmetals can be very exothermic:
2H2(g) + 2O2(g)  2H2O(l)
∆Ho = –571.7 kJ
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57. Hydrogen
It can either gain another electron to form the hydride ion, H–, or lose its electron to become H+:
2Na(s) + H2(g)  2NaH(s)
2H2(g) + O2(g)  2H2O(l)
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58. Hydrogen
H+ is a proton.
The aqueous chemistry of hydrogen is dominated by H+(aq).
(acids – next semester)
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59. Group 6A Oxygen, sulfur, and selenium are nonmetals.
Tellurium is a metalloid.
The radioactive polonium is a metal.
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60. Oxygen There are two allotropes of oxygen: There can be three anions:
O3, ozone
There can be three anions:
O2−, oxide
O22−, peroxide
O21−, superoxide
It tends to take electrons from other elements (oxidation).
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61. Sulfur Sulfur is a weaker oxidizer than oxygen.
The most stable allotrope is S8, a ringed molecule.
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62. Group 7A: Halogens The halogens are prototypical nonmetals.
The name comes from the Greek words halos and gennao: “salt formers”.
All halogens are diatomic: X2.
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63. Group 7A: Halogens
The chemistry of the halogens is dominated by gaining an electron to form an anion:
X2 + 2e–  2X–
Fluorine is one of the most reactive substances known:
2F2(g) + 2H2O(l)  4HF(aq) + O2(g) DH = –758.9 kJ
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64. Group 7A: Halogens They have large, negative electron affinities.
Therefore, they tend to oxidize other elements easily.
They react directly with metals to form metal halides.
Chlorine is added to water supplies to serve as a disinfectant
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65. Group 7A: Halogens
Chlorine is added to water supplies to serve as a disinfectant
e.g. swimming pools:
Cl2(g)+ H2O(l)  HCl(aq) + HOCl(aq)
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66. Group 8A: Noble Gases
The noble gases have astronomical ionization energies.
Their electron affinities are positive.
Therefore, they are relatively unreactive.
They are found as monatomic gases.
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67. Group 8A: Noble Gases Xe forms three compounds:
XeF4 (at right)
XeF6
Kr forms only one stable compound:
KrF2
The unstable HArF was synthesized in 2000.
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68. Sample Integrative Exercise 7 (p. 285)
The element bismuth (Bi, atomic number 83) is the heaviest member of group 5A. A salt of the element, bismuth subsalicylate, is the active ingredient in Pepto-Bismol, an over-the-counter medication for gastric distress.
The covalent atomic radii of thallium (Tl) and lead (Pb) are 1.48 Å and 1.47 Å, respectively. Using these values and what you know about trends in atomic radii, predict the covalent atomic radius of the element bismuth. Explain your answer.
What accounts for the general increase in atomic radius going down the group 5A elements?
3. Another major use of bismuth has been as an ingredient in low-melting metal alloys, such as those used in fire sprinkler systems and in typesetting. The element itself is a brittle white crystalline solid.
How do these characteristics fit with the fact that bismuth is in the same periodic group with such nonmetallic elements as nitrogen and phosphorus?
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69. Sample Integrative Exercise 7 (p. 285)
Bi2O3 is a basic oxide. Write a balanced equation for its reaction with dilute nitric acid. If 6.77 g of Bi2O3 is dissolved in dilute acidic solution to make up 500 mL of solution, what is the molarity of the Bi3+ ion?
209Bi is the heaviest stable isotope of any element. How many protons and neutrons are present in this nucleus?
6. The density of Bi at 25oC is g/cm3. How many Bi atoms are present in a cube of the element that is 5.00 cm on each edge? How many moles of the element are present?
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