1. Chapter 7- Periodic Properties of the Elements

2. Periodic table- One of the most significant tools chemists have for organizing and remembering chemical facts
In Chapter 6- what did we say was primarily responsible for the properties/reactivities of elements?

3. Electron configuration- repeating pattern on periodic table
Responsible for periodic nature of the elements
Elements in same column (group)= same number of valence electrons
What were valence electrons again?
Elements in same group- what else is special about them?

4. Why do elements in the same group have similar properties?

5. These differences can also be attributed to electron configuration
What you haven’t heard- elements in same group share properties, but also have differences
Ex: Oxygen vs. Sulfur- both in group 6A- bond/react similarly
Oxygen: colorless gas at room temp
Sulfur: yellow solid at room temp
These differences can also be attributed to electron configuration
What shell (value of n) are the valence electrons of O in?
What shell are the valence electrons of S in?

6. Sec. 7.1
Mendeleev- created one of the first widely accepted periodic table
Inspired by his favorite card game, solitaire, Mendeleev created a set of cards of the elements- one card for each element currently discovered.
Each element card had information on it: the element’s name, atomic mass, properties, and compounds they formed with other elements

7. Mendeleev organized his periodic table in order of increasing atomic mass- he spread all the cards out, placed them in a big line according to increasing mass
When ordered this way, he noticed a pattern arise among the elements- similar properties actually repeated themselves in a regular pattern (a periodic pattern)

8. Periodic Table Build!
You are given a set of element cards- all info on these cards should be used in order to organize the elements you are given into the first periodic table (correctly)

9. Discovery of chemical elements has been an ongoing process since ancient times
Though many elements are stable- tend to form compounds
Not many elements found in elemental form naturally- why’s that?
Many elements went undiscovered due to the fact that they naturally form compounds

10. Early 1800s- advances in chem allow for isolation of elements from their compounds
These advances= many elemental discoveries
With number of elements increasing, the demand for a method/way of organizing them increased
Many scientists tried many different classifying schemes
In steps Mendeleev with his crazy table, rambling on about how his way of organizing the elements was the best way
How did he organize it?

11. Why did you end up leaving holes in your periodic table?
What did you end up doing with the holes that were left?

12. Mendeleev ended up with the same holes in his periodic table- he insisted that elements with similar properties go in the same families- this forced him to leave holes in his table
Using his awesome table, Mendeleev boldly went were no man had gone before and decided he would predict the existence of these undiscovered elements- he went as far as predicting their atomic mass, melting point, color, density, specific heat, boiling point, and the compounds it would form with oxygen and chlorine

13. He called the element hole under Aluminum “eka- Aluminum” and the element hole under Silicon “eka- Silicon”
When Gallium and Germanium were discovered, they very closely matched the properties Mendeleev has predicted
This proved how useful Mendeleev’s periodic table was- it can be used to predict properties of undiscovered elements

14. Property
Mendeleev’s prediction for eka-Silicon (1871)
Germanium properties (1886)
Atomic weight 72
72.59
Density
5.5 g/mL
5.35 g/mL
Specific heat
0.305 (J/gK)
0.309 (J/gK)
Melting Point
High
947 °C
Color
Dark gray
Grayish white
Oxide Formula
XO2
GeO2
Chloride Formula
XCl4
GeCl4
Boiling Point of chloride compound
A little under 100 °C
84 °C

15. Even though Mendeleev’s table was the bees knees, it had its problems
Do you remember which elements broke the trend of increasing atomic mass? Why did you place them out of order?

16. After Rutherford’s model of the atom, a physicist named Henry Mosely decided he would shoot high energy electrons at atoms
These atoms emitted X-rays of specific and unique frequencies- frequency emitted generally increased as atomic mass increased
Assigned each element a specific whole number values based on the frequency of X-ray emitted- atomic number

17. Stated atomic number= number of protons in nucleus
Ordered periodic table in order of increasing atomic number- this is how the modern periodic table is arranged
Ordering in increasing atomic number- clarified the problems seen in Mendeleev’s table
This new arrangement also allowed for the identification of holes/undiscovered elements- property predictions as well- led to discovery of even more elements

18. Sec. 7.2- Effective Nuclear Charge

19. Electrons- negative charge Nuclei- contain protons- positive charge
Electron attracted to nucleus- atom properties depend on electron configurations and how strongly their outer electrons are attracted to the nucleus
Coulomb’s Law= strength of interaction between 2 electrical charges depends on magnitudes of charge and distance between them
Attractive force between electrons and nucleus increases as nuclear charge increases
Decreases as electrons move further away

20. Many electron atoms- all electrons attracted to nucleus and repelled by one another at the same time
Due to number of repulsions, it is difficult to measure all the forces
However, we can estimate the net attraction of each electron to the nucleus- must consider how single electron interacts with “average” environment created by nucleus and other electrons

21. This allows us to treat each electron as if it were moving in a net electric field created by nucleus and electron density of other electrons
Effective Nuclear charge (Zeff)- the nuclear charge felt by an electron- takes into consideration the “shielding” effects the inner shell electrons have; this is the “net” nuclear charge felt by individual electron
Shielding- caused by the repulsions between electrons
Effective nuclear charge will always be smaller than the actual nuclear charge

22. Valence electron- attracted to nucleus in atom, but also repelled by inner shell electrons- the repulsions from inner shell electron partially “cancel out” some of the attraction felt from the nucleus (shielding)
Zeff = Z – S
Zeff= effective nuclear charge, Z= nuclear charge (# protons in nucleus), S= shielding constant

23. S is portion of nuclear charge that is “cancelled out” by repulsion of inner shell electrons
S is normally close to the value of inner shell electrons
Electrons in same valence shell are not good at shielding one another

24. Zeff For a Sodium Atom

25. Why do you think this is? What changes as we move from left to right?
Periodic trend: Zeff for valence electrons increases as we move from left to right across a period
Why do you think this is? What changes as we move from left to right?
Going down a group, Zeff slightly increases, but not by much- not nearly as much as it does moving across a period
Would expect the Zeff to be about the same for elements in same group- why?

26. Sec. 7.3- Sizes of Atoms and Ions

27. Size of atoms and ions is a somewhat strange concept, but is also extremely important
Atoms- not hard, spherical objects- if you remember from chapter 6, electron clouds are not really “hard” boundaries- no sharp cut off for the boundary at which electron distribution becomes zero
We define atomic size in several ways- mostly based on the distance between atoms in various situations

28. Argon atoms- when these gas particle collide, they easily bounce off one another without reacting/interacting all that much
This is due to the fact that the electron clouds of these gas atoms will not mesh together to any significant extent – no attraction because these atoms are not charged, and the clouds will actually somewhat repel one another

29. When atoms collide in this way, the collision is sometimes referred to as a “non-bonding” collision- since the atoms are simply ricocheting off one another
Nonbonding atomic radius- distance between two nuclei during a “non-bonding” collision

30. When atoms are bonded to one another- like Cl2- there is some sort of attractive force between the bonded atoms- hence why they are bonded
Due to this attractive force between bonded atoms, the nuclei of the bonded atoms will be closer to one another than they would be if they were not bonded- the bond kind of squeezes their electron clouds closer together

31. Bonding atomic radius- the distance between the nuclei of bonded atoms
How does bonding atomic radius compare to nonbonding atomic radius?

32. Through observation,/experiments, scientists have gathered data on the bonding atomic radius for every element- each element assigned its own bonding radius value
With these values, we are able to tell how long a bond will be between any combination of bonded atoms in molecules
Ex: For C, bonding atomic radius= 0.77 Å, and Cl= 0.99 Å
In the compound CCl4, we can estimate the bond length of a C-Cl bond to be 1.76 Å (In reality, the bond length is 1.77 Å, which is pretty close to the estimated value)

33. Periodic Trends in Atomic Radii
Two trends for atomic radius:
1- Atomic radius tend to increase from top to bottom (going down) a group. Why do you think this is? What happens to the value of n as we go down a group?
2- Atomic radius tends to decrease when moving from left to right across a period. Major influence for this is the Zeff. Why does Zeff have an effect on atomic radius?

34. Periodic Trends in Ionic Radii
Ionic radii- determined by distance between nuclei of ions
Size of ion depends on effective nuclear charge, number of electrons it possesses, and the orbitals in which its valence electrons are located
When a cation forms (positive ion), does this mean that electrons were gained or lost?
When an anion forms (negative ion), does this mean that electrons were gained or lost?

35. If cations form when electrons are lost, how do you think the ionic radius of a cation compares to that of the regular atomic radius?
If anions form when electrons are gained, how does the ionic radius of an anion compare to that of the regular atomic radius?

36. Ionic radius of cation < atomic radius
Ionic radius of anion > atomic radius
Sooo… cations are smaller than their “parent” atoms, and anions are larger than their “parent” atoms
Also, for ions with same charge, size increases as we move down a group

37. Based on periodic trends for atomic radius and ionic radius, order to following atoms and ions in order of increasing size (smallest to largest): Mg+2, Ca+2, and Ca

38. Isoelectronic series: groups of ions all containing same number of electrons
Ex: O2-, F-, Na+, Mg2+, Al3+ … all have how many electrons
If members of isoelectronic series are listed in order of increasing atomic number, effective nuclear charge will increase. Why?
If nuclear charge increases, what will happen to the ionic radius?
List the ions above in order of increasing ionic radius

39. Sec. 7.4- Ionization Energy

40. How easily electrons can be removed from atoms/ions has major impacts on chemical behavior
Ionization energy- minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion
First ionization energy (I1)- energy required to remove the first electron from a neutral atom

41. Second Ionization energy (I2)- energy needed to remove the second electron
This pattern continues for third, fourth, fifth ionization energies (continues until atom is out of electrons)
Which ionization energy do you think is the smallest? Why?

42. First ionization energy for the sodium atom is the energy required for the process:
I2 for sodium is energy associated with this process:
Ionization energies- always positive, due to the fact that energy must be applied to the atom to pull an electron away

43. First ionization energy is always the smallest- the first electron is easiest to pull away- partially due to the fact that the atom is neutrally charged at this point
Higher ionization energies= more difficult to remove electrons

44. This is due to the fact that after the first ionization, the atom becomes positively charged- creating a stronger attraction between the electrons and the nucleus
There is also a trend when it comes to inner shell vs. valence shell electrons- which electrons do you think would be easiest to remove?

45. There is a sharp increase in ionization energy once all valence electrons have been removed- greater increase than from the first to second ionization energy (if more than 1 valence electron)
Inner shell electrons- much more difficult to remove than valence electrons- why?

46. The fact that valence electrons have such low ionization energies than the inner shell electrons supports the fact that valence electrons are mostly responsible for how something bonds/ chemically reacts
The valence electrons will be the only ones willing to transfer/share- form ionic/covalent bonds
Inner electrons too tightly bound to nucleus to be lost from atom or even be shared with another atom

47. Predict which atom will have the largest second ionization energy: Calcium, Sodium, and Sulfur

48. Periodic Trend of First Ionization Energy
First ionization energy, I1, typically increases going left to right across the same period (some exceptions to this rule)
Ionization energy typically decreases going down a group- as atomic number increases

49. The s and p block element show a larger range in values of I1 than the d block (transition metals) do. Generally, the ionization energy of the transition metals increases slowly as we move from the left to the right across a given period.
In general, smaller atoms have higher ionization energies- why do you think this happens?

50. The energy needed to remove an electron depends on:
Effective nuclear charge
Distance of the electron from the nucleus
Increasing effective nuclear charge or decreasing the distance between the electron and nucleus will increase the attraction between the electron and the nucleus
As attraction to nucleus increases, it becomes more difficult to remove the electron- more attraction = more required energy= higher ionization energy

51. Moving left to right across a period, what happens to:
Effective nuclear charge?
Atomic size?
So energy needed to remove electron will:
When moving down a group, what happens to:
So, energy needed to pull electron away will:

52. Arrange the following from smallest to largest first ionization energy: Ne, Na, P, Ar, K

53. Sec. 7.5- Electron Affinities

54. Measures the ease in which an atom will gain an electron
Electron affinity- energy change that occurs when an electron is added to a gaseous atom
Measures attraction (affinity) of an atom to an added electron
Measures the ease in which an atom will gain an electron

55. Electron affinities are typically negative energy values (this indicates that energy is released when electron added)
If the electron affinity value is positive, this means the atom will not gain an electron
The more negative the electron affinity is, the more the atom will want to gain an electron

56. No real periodic trend when it come to electron affinities- however, some groups of elements have characteristically low/high affinities
The Halogens- Have the most negative electron affinity values… why?
The Noble Gases- do you think their electron affinities will be negative or positive? Why?

57. Sec. 7.6- Metals, Nonmetals, and Metalloids

58. Elements can be broadly grouped into the categories of metals, nonmetals, and metalloids
Metalloid staircase- marked on your periodic tables
Left of staircase- metals (any exceptions?)
Right of staircase- Nonmetals
Touching staircase on two sides- metalloids (Boron, silicon, germanium, arsenic, antimony, tellurium)- form a border between metals and nonmetals
What category do most elements on the periodic table belong to?

60. Metallic character- a measure of how much an element exhibits the physical and chemical properties of metals
Metallic character- Increases moving down a group; decreases moving left to right across a period

61. Metals Shiny luster; various colors- often silvery
Good conductors of heat and electricity (this is due to their metallic bonding)
Most are Malleable- can be pounded into thin sheets
Most are Ductile- can be pulled into a thin wire

62. High melting points- all are solid at room temp except one… which one?
Tend to have low ionization energies- what does this mean they tend to do with their valence electrons?
Tend to form what charges?

63. When metals undergo chemical changes, they typically oxidize (Oxidation is Loss) due to low ionization energies/ electron loss
For s block metals- they will lose electrons in the outermost s shell- forming positive charges- either 1+ or 2+ charges form

64. For p block metals- They will either lose just the outer p electrons, or both the outer s and p electrons
Al3+ - what block did the electrons this metal lost belong to?
Sn2+ - what block did the electrons this metal lost belong to?
Sn4+- what block did the electrons this metal lost belong to?
Charges formed by transition metals do not follow a distinctive pattern- often able to form more than one positive ion

65. When metals bond with nonmetals, they typically form an ionic compound
When metals left outside, typically combine with the nonmetal oxygen (abundant in our air) to form an ionic compound:
Ni O2 →
Which one of these reactants- Ni or O2, will be easier to pull electrons away from? How do you know this?
Ionic bonding- one atom will lose electrons- the other will gain- which element is losing electrons in this reaction? Which one is gaining?

66. Whenever metals react with oxygen, we call the products metal oxides- metals oxides are fairly common ionic compounds due to the abundance of oxygen in nature
Metal oxides have a special characteristic- most of them are basic when in solution
When metal oxides are able to dissolve in water (if soluble) they react to form metal hydroxides:
Metal oxide + water → metal hydroxide

67. If you look at the net ionic equation, the basicity of metal oxides is due to the oxide ion, which reacts with water as follows:

68. Metal oxides also display their basicity when reacting with acids:
Metal oxide + acid → salt + water
What kind of reaction does that look like?

69. Nonmetals Nonmetals- vary greatly in their physical appearance
No luster, and are poor conductors of heat and electricity
Melting points- typically lower than those of metals (Notice that all gases on the periodic table are nonmetals)
Not malleable/ductile- may be very hard, brittle, or soft

70. Nonmetals- how do their ionization energies compare to those of metals
Nonmetals- how do their ionization energies compare to those of metals? What about their electron affinities?
So when nonmetals bond with metals, what will they typically do?
Ex: Aluminum reacting with bromine

71. When nonmetals bond ionically (with other metals) they will gain enough electrons to fill their outermost occupied p subshell
When nonmetals bond with other nonmetals- this forms a covalent bond (electrons are shared)
Nonmetal oxides- acidic!
When dissolved in water, nonmetal oxides react to form an acid
Nonmetal oxide + water → acid
Ex:

72. Nonmetal oxides will also react with bases to form salt and water
Nonmetal oxide + base → salt + water
Ex:

73. Metalloids
Metalloids- kind of like the strange child of metal/nonmetal
Metalloids have properties that are both metallic and nonmetallic- depends on which metalloid you are dealing with
Ex: Silicon- looks like a metal, but when you hit it with a hammer, it shatters (not malleable like metal)

74. Several metalloids- most notably silicon- are semiconductors and are used in computer chips/circuits
Pure Silicon= electrical insulator… but if you add specific impurities, its electrical conductivity dramatically increases- this makes it possible to control electrical conductivity by simply controlling chemical composition

75. Sec. 7.7- Group Trends for the Active Metals

76. We have seen that elements in the same group posses general similarities
There are trends found without groups- using periodic table and our knowledge of electron configuration to examine the chemistry of the alkali metals, and the alkaline earth metals

77. Group 1A- Alkali Metals Alkali Metals= metals in group 1A
Soft, metallic solids, silvery, metallic luster, and good conductors of heat/electricity
They have a fairly low density and melting point (compared to other metals)
Moving down the alkali metals: melting point decreases, density increases, what about atomic radius? First ionization energy?

78. For each row (period) on the periodic table, the alkali metals have the lowest I1 value- what does this reflect about its outer s electron?
Due to the low ionization energy- alkali metals lose electrons very easily= very very very reactive
All will form what charge?

79. Will combine directly with most nonmetals
Alkali metals only exist as compounds- partially due to high reactivity
Will combine directly with most nonmetals
Alkali metals will combine with hydrogen to form hydrides and sulfur to produce sulfides:
Na + H2 →
Na + S →
These reactions happen with all alkali metals
When reacting with H2, alkali metals will give up their electron to H, which will then form a hydride ion:

80. All alkali metals are violently reactive with water
When they combine with water, they produce hydrogen gas and a metal hydroxide

81. The reaction between alkali metals and water is extremely exothermic- which means it releases a lot of heat- most of the time this is enough heat to ignite the H2 gas that is also produced… this makes a boom
How do you think the violence of this reaction changes as you go down a group?

82. Alkali metal ions= colorless in compounds/ solutions- But, when placed in flame, they all emit a characteristic color of light
When placed in flame, the alkali metal ions are reduced to gaseous metal atoms, and the valence electron absorbs energy from the heat of the flame
What happens when the single valence electron of an alkali metal absorbs energy?

83. When electrons become excited, they will want to immediately return to ground state
What will the excited valence electron have to do in order to return to ground state?

84. When releasing energy, the excited electron releases a photon with unique energy- due to this unique energy, the photon emitted will have a special wavelength- this will correspond with a specific color of light
Every metal ion releases a unique color of light when placed in a flame- why?
Ex: Sodium will release yellow light after being excited- valence electron in 3s jumps to 3p when excited- when returning back to ground state (3p to 3s) it releases a photon with a wavelength of 589 nm (yellow light)

85. Group 2A: The Alkaline Earth Metals

86. Alkaline earth metals are harder than alkali metals, have higher melting points, and are more dense
Going down the alkaline earth metals group, melting points decrease, density increases, atomic radius? First ionization energy?
How do you think their reactivity compares to the alkali metals?
What about reactivity going down the alkaline metal group?

87. The way the alkaline earth metals react with water shows the trend of reactivity (increases going down the alkaline earth metals)- none are as reactive as the alkali metals
Beryllium does not react with water, but every other alkaline earth metal does
Ex: Magnesium and water (all other alkaline earths- besides beryllium- react with water the same way)

88. Which electrons will the alkaline earth metals lose?
What ion will all alkaline earth metals form?
Alkaline earth metal with chlorine and oxygen:
Heavier alkaline earths (lower in group) are even more reactive to nonmetals

89. Alkaline earth metals are similar to alkali metals when in the presence of a flame
When heated to very high temperatures, their valence electrons will become excited- once excited they will release photons with a characteristic color (energy/wavelength) to return to ground state
Flame test analysis can help us identify the type od metal in an unknown compound

90. Sec. 7.8- Group Trends of Selected Nonmetals

91. Hydrogen
Very strange- doesn’t know exactly what to do with its valence electron- will be happy losing it, or gaining 1 more
Hydrogen does not lose its valence electron as easily as the alkali metals
When bonding with other nonmetals, hydrogen will share its valence electron instead of losing it
When reacting with alkali metals, will gain an electron to form H- (hydride ion)
In aqueous chemistry, hydrogen loses its electron to form H+, hydrogen cation (acid!)

92. Group 6A- The Oxygen Family
Not many shared physical characteristics- mix of nonmetals, one metalloid, and a metal
Oxygen- gas at room temp, the rest of the group are solids
Oxygen, sulfur, selenium- typical nonmetals; tellurium- metalloid; Polonium- radioactive metal
All have 2 valence electrons and will typically form what type of charge:

93. Oxygen- found in 2 naturally occurring molecular forms
O2= dioxide (typically referred to as just “oxygen”), and O3 = ozone
O2 and O3 are allotropes- different forms of the same element in the same state (O2 and O3 are gases)
O2 is necessary for life- we use it for respiration

94. O3 is found in upper atmosphere and polluted air- damaging to health, but also necessary for life on Earth- ozone absorbs a large part of the UV radiation emitted from the Sun
O2 has a great tendency to attract electrons (oxidizes other elements)- when combining with metal, forms oxide ion O2-
Sulfur also has a tendency to gain electrons from other elements to form sulfides containing the S2- ion (also does this in combination with metals)

95. Group 7A- The Halogens
Group 7A= the Halogens (group beginning with fluorine)
All are typical nonmetals (astatine is rare and radioactive, so not much data)
All (except astatine) are diatomic
Going down the halogen group- what do you think happens to melting point?

96. All have different colored vapors- F2= pale yellow, Cl2= yellow-green, Br2= reddish-brown, I2= violet vapor
Highly negative electron affinities- meaning what?

97. The halogens are the most reactive nonmetals due to their extremely negative electron affinities
What ion will all halogens form?
Fluorine- most reactive, removes electrons from almost any substance it comes into contact with

98. Halogens will react with metals to form metal halides:
Halogens also react with hydrogen to form gaseous hydrogen halide compounds:

99. Group 8A- The Noble Gases
Group 8A= The Noble Gases- far right of the periodic table
All are nonmetals and gases at room temperature
All are monatomic- single atom, not a molecule (not diatomic)
Completely filled s and p subshells- this makes them happy (full valence shells)

100. All have very large ionization energies- meaning what?
All have positive electron affinities- meaning what?
Stable electron configuration= very unreactive
Do not readily form compounds
Radon- radioactive