1. Chapter 7 Periodic Properties of the Elements

2. Development of the Periodic Table
Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped.

3. Mendeleev and the Periodic Table
Chemists mostly credit Mendeleev because he also used chemical properties to organize the table and predicted some missing elements and their expected properties, including germanium.

4. Properties and Electron Configuration
The properties of the elements follow a periodic pattern.
Elements in the same column have similar properties.
The elements in a period show a pattern that repeats.
The quantum-mechanical model explains this because the number of valence electrons and the types of orbitals they occupy are also periodic.

5. The Noble Gas Electron Configuration: ns2np6
The noble gases have eight valence electrons.
Except for He, which has only two electrons
They are especially nonreactive.
He and Ne are practically inert.
The reason the noble gases are so nonreactive is that the electron configuration of the noble gases is especially stable: full valence level.

6. The Alkali Metals
The alkali metals have one more electron than the previous noble gas.
In their reactions, the alkali metals tend to lose one electron, resulting in the same electron configuration as a noble gas.
Forming a cation with a 1+ charge

7. The Halogens
The electron configurations of the halogens all have one fewer electron than the next noble gas.
In their reactions with metals, the halogens tend to gain an electron and attain the electron configuration of the next noble gas, forming an anion with charge 1−.
In their reactions with nonmetals, they tend to share electrons with the other nonmetal so that each attains the electron configuration of a noble gas.

8. Eight Valence Electrons
Quantum mechanical calculations show that eight valence electrons should result in a very unreactive atom.
an atom that is very stable
the noble gases have eight valence electrons and are all very stable and unreactive.
He has two valence electrons, but that fills its valence shell.
Conversely, elements that have either one more or one less electron should be very reactive.
the halogen atoms have seven valence electrons and are the most reactive nonmetals.
the alkali metals have one more electron than a noble gas atom and are the most reactive group of metals.

9. Electron Configuration & Ion Charge2 Ne 3 Ar 4 5 6
Ar = 1s22s22p63s23p6
Cl = 1s22s22p63s23p5
Cl1− anion = 1s22s22p63s23p6
S = 1s22s22p63s23p4
S2− anion = 1s22s22p63s23p6
K = 1s22s22p63s23p6 4s1
K1+ cation = 1s22s22p63s23p6
Ca = 1s22s22p63s23p6 4s2
Ca2+ cation = 1s22s22p63s23p6

10. Atomic Number
Mendeleev’s table was based on atomic masses. It was the most fundamental property of elements known at the time.
About 35 years later, the nuclear atom was discovered by Ernest Rutherford.
Henry Moseley developed the concept of atomic number experimentally. The number of protons was considered the basis for the periodic property of elements.

11. Periodicity
Periodicity is the repetitive pattern of a property for elements based on atomic number.
The following properties are discussed in this chapter:
Sizes of atoms and ions
Ionization energy
Electron affinity
Some group chemical property trends
First, we will discuss a fundamental property that leads to many of the trends, effective nuclear charge.

12. Effective Nuclear Charge
Many properties depend on attractions between valence electrons and the nucleus.
Electrons are both attracted to the nucleus and repelled by other electrons.
The forces an electron experiences depend on both factors.

13. Shielding
Outer valence electrons are shielded from the nucleus by the core electrons.
Screening or shielding effect
Outer electrons do not effectively screen for each other.
The shielding causes the outer electrons to not experience the full strength of the nuclear charge.

14. Effective Nuclear Charge
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.
Effective nuclear charge is a periodic property:
It increases across a period.
It increases slightly down a group.
Although the number of core electrons stays the same across a period, the number of protons increases which accounts for the effective nuclear charge increasing across a period.
The effective nuclear charge increases slightly as we go down a column because the more diffuse core electron cloud is less able to screen the valence electrons from the nuclear charge.

15. Shielding & Effective Nuclear Charge
Valence (2s1) electron
Valence (2s2) electron
Nucleus (4+)
Valence (2s2) electron
Lithium
Effective Nuclear Charge
Zeff  (3+) – 2
 1
Beryllium
Effective Nuclear Charge
Zeff  (4+) – 2
 2

16. Trend in Atomic Radius – Main Group
There are several methods for measuring the radius of an atom, and they give slightly different numbers.
van der Waals radius = nonbonding
covalent radius = bonding radius
atomic radius is an average radius of an atom based on measuring large numbers of elements and compounds
Atomic radius increases down a group.
valence shell farther from nucleus
effective nuclear charge constant
Atomic radius decreases across period (left to right).
adding electrons to same valence shell
effective nuclear charge increases
valence shell held closer

17. Periodic Table Trends – Atomic Size
Atomic Radius Decreases
Atomic Radius Increases

18. Choose the Larger Atom in Each Pair
N or F?
C or Ge?
N or Al?
Al or Ge?
N or F? N farther left
C or Ge? Ge farther down
N or Al? Al farther left and down
Al or Ge? opposing trends

19. Sizes of Ions Determined by interatomic distances in ionic compounds
Ionic size depends on
the nuclear charge.
the number of electrons.
the orbitals in which electrons reside.

20. Sizes of Ions Cations are smaller than their parent atoms:
The outermost electron is removed and repulsions between electrons are reduced.
Anions are larger than their parent atoms:
Electrons are added and repulsions between electrons are increased.

21. Size of Ions— Isoelectronic Series
In an isoelectronic series, ions have the same number of electrons.
Ionic size decreases with an increasing nuclear charge.
An isoelectronic series (10 electrons)
Note increasing nuclear charge with decreasing ionic radius as atomic number increases

22. Which ion is the smallest?
Mg2+
Ca2+

23. Ionization Energy (I)
The ionization energy is the minimum 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 the energy required to remove the second electron.
Note: The higher the ionization energy, the more difficult it is to remove an electron!

24. Ionization Energy
It requires more energy to remove each successive electron.
When all valence electrons have been removed, it takes a great deal more energy to remove the next electron (a core electron).

25. Periodic Trends in First Ionization Energy (I1)
I1 generally increases across a period.
I1 generally decreases down a group.
The s- and p-block elements show a larger range of values for I1. (The d-block generally increases slowly across the period; the f-block elements show only small variations.)
Going across a period the effective nuclear charge increases which makes it harder to pull off an electron.
Going down a group the electrons are pulled from a shell which is furthest from the nucleus where they are held less tightly. Thus the ionization energies decrease down a group.
All ionization energies for atoms are positive: energy must be absorbed to remove an electron.

26. Periodic Table Trends – Ionization Energy
Ionization Energy Increases
Ionization Energy decreases

27. Choose the atom in each pair with the larger first ionization energy
Al or S
As or Sb
N or Si
O or Cl?
Al or S, sulfur has >Zeff
As or Sb, As e- held more tightly
N or Si, N is farther up & right
O or Cl opposing trends

28. Consider the following successive ionization energies (kJ/mol):
IE IE IE IE IE IE IE7
, ,345
Which element in period 3 would most likely show this trend in ionization energies? Mg Al Si P S

29. Electron Configurations of Ions
Cations: The electrons are lost from the highest energy level (n value).
Li+ is 1s2 (losing a 2s electron).
Fe2+ is 1s22s22p63s23p63d6 (losing two 4s electrons).
Anions: The electron configurations are filled to ns2np6; for example, F– is 1s22s22p6 (gaining one electron in 2p).
It may seem odd that 4s electrons are removed before 3d electrons in forming transition-metal cations. After all, in writing electron configurations, we added the 4s electrons before the 3d ones. In writing electron configurations for atoms, however, we are going through an imaginary process in which we move through the periodic table from one element to another. In doing so, we are adding both an electron to an orbital and a proton to the nucleus to change the identity of the element. In ionization, we do not reverse this process because no protons are being removed.
Electrons added to an atom to form an anion are added to the empty or partially filled orbital having the lowest value of n.

30. Electron Affinity
Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom:
Cl + e– Cl–
It is typically exothermic, so, for most elements, it is negative!
The greater the attraction between an atom and an added electron, the more negative the atom’s electron affinity.

31. The trends in electron affinity are not as evident as they are for ionization energy.
The halogens, which are one electron shy of a filled p subshell, have the most negative electron affinities. By gaining an electron, a halogen atom forms a stable anion that has a noble-gas configuration.
The addition of an electron to a noble gas, however, requires that the electron reside in a higher-energy subshell that is empty in the atom. Because this is unfavorable the EA is highly positive.
Electron affinities do not change greatly as we move down a group. Going down the halogen group, however, the added electron is being placed in a subshell that is further from the nucleus. The electron-nucleus attraction decreases because of this increased distance. But the electron-electron repulsions are similarly reduced so there is not a clear trend.

32. Periodic Table Trends – Electron Affinity
Electron Affinity Increases
Electron Affinity decreases

33. Metal, Nonmetals, and Metalloids
Decreasing metallic character

34. Metals Differ from Nonmetals
Metals tend to form cations.
Nonmetals tend to form anions.
Note the special property of hydrogen.

35. Metals Most of the elements in nature are metals.
Properties of metals:
Shiny luster
Conduct heat and electricity
Malleable and ductile
Solids at room temperature (except mercury)
Low ionization energies/form cations easily

36. Metal Chemistry
Compounds formed between metals and nonmetals tend to be ionic.

37. Nonmetals
Nonmetals are found on the right hand side of the periodic table.
Properties of nonmetals include the following:
Solid, liquid, or gas (depends on element)
Solids are dull, brittle, poor conductors
Large negative electron affinity, so they form anions readily

38. Nonmetal Chemistry
Carbon dioxide added to water dissolves in the water and combines with water to form carbonic acid which turns the water slightly acidic as can be seen with this bromthymol blue indicator.
Substances containing only nonmetals are molecular compounds.

39. Metalloids
Metalloids have some characteristics of metals and some of nonmetals.
Several metalloids are electrical semiconductors (computer chips).

40. Alkali Metals Alkali metals are soft, metallic solids.
They are found only in compounds in nature, not in their elemental forms.
Typical metallic properties (luster, conductivity) are seen in them.

41. Alkali Metal Properties
They have low densities and melting points.
They also have low ionization energies.

42. Alkali Metal Chemistry
Their reactions with water are famously exothermic.

43. Flame Tests
Qualitative tests for alkali metals include their characteristic colors in flames.
These are caused by electronic transitions.

44. 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.
They readily form +2 cations, losing the 2 valence electrons.

45. Alkaline Earth Metals
Beryllium does not react with water, and magnesium reacts only with steam, but the other alkaline earth metals react readily with water.
Reactivity tends to increase as you go down the group.

46. Group 6A—Increasing in Metallic Character down the Group
Oxygen, sulfur, and selenium are nonmetals.
Tellurium is a metalloid.
The radioactive polonium is a metal.

47. Group 7A—Halogens The halogens are typical nonmetals.
They have highly negative electron affinities, so they exist as anions in nature.
They react directly with metals to form metal halides.

48. Group 8A—Noble Gases
The noble gases have very large ionization energies.
Their electron affinities are positive (can’t form stable anions).
Therefore, they are relatively unreactive.
They are found as monatomic gases.