1. PRINCIPLES OF CHEMISTRY I CHEM 1211
CHAPTER 8
DR. AUGUSTINE OFORI AGYEMAN
Assistant professor of chemistry
Department of natural sciences
Clayton state university

2. CHAPTER 8
PERIODIC PROPERTIES OF
THE ELEMENTS

3. CLASSIFICATION OF THE ELEMENTS
- Elements in a given group have similar chemical properties
because the outer-shell electron arrangements are similar
Group IIA elements
Be: 1s22s2
Mg: 1s22s22p63s2
Ca: 1s22s22p63s23p64s2
Sr: 1s22s22p63s23p64s23d104p65s2

4. CLASSIFICATION OF THE ELEMENTS
- The last electron in an element’s electron configuration
causes the difference in the electron configuration of the
preceding element and is referred to as the
distinguishing electron
Homework
Write notes (one page) on the different classifications of the
elements based on electronic properties. Briefly describe the
s-area, p-area, d-area, and the f-area.

5. CLASSIFICATION OF THE ELEMENTS
Elements can be classified as
Metals or Nonmetals
Based on physical properties
Elements can also be classified as
Noble-gas, Representative,
Transition, or Inner Transition
- Based on electron configuration

6. CLASSIFICATION OF THE ELEMENTS
Noble-gas Elements
- Group VIIIA (18) elements on the periodic table
(far right column)
- Gases at room temperature
- Little tendency to form chemical compounds
- Electron configuration ends in p6
- Completes p subshell (except Helium)
- Nonmetals

7. Representative Elements
CLASSIFICATION OF THE ELEMENTS
Representative Elements
- Elements in the
s-area (Groups IA and IIA)
first five columns of the p-area
(Groups IIIA, IVA, VA, VIA, and VIIA)
- Metals and nonmetals

8. CLASSIFICATION OF THE ELEMENTS
Transition Elements
- Elements in the d-area of the periodic table
- Groups IIIB (3), IVB (4), VB (5), VIB (6), VIIB (7),
VIIIB (8, 9, 10), IB (11), and IIB (12)
- Distinguishing electron in a d subshell
- Metals

9. Inner Transition Elements
CLASSIFICATION OF THE ELEMENTS
Inner Transition Elements
- Elements in the f-area of the periodic table
- The two-row block of elements below the main table
- Distinguishing electron in an f subshell
- Metals

10. VALENCE ELECTRONS CORE ELECTRONS
- Electrons in the highest principal quantum number
of an atom, and any electrons in an unfilled subshell
from a lower shell
CORE ELECTRONS
- Electrons in filled subshells that have lower
principal quantum numbers

11. VALENCE ELECTRONS
- Representative elements in the same group of the periodic table
have the same number of valence electrons
- The number of valence electrons for representative elements is the
same as the group number (with A) in the periodic table
- The maximum number of valence electrons for any given element
is eight

12. VALENCE ELECTRONS
- Not all electrons in a given atom participate in bonding
- Only valence electrons are available for bonding
- For representative and noble-gas elements
these electrons are always found in the s or p subshells

13. VALENCE ELECTRONS
- The electron configuration can be used to determine
the number of valence electrons of an atom
C: 1s22s22p2
O: 1s22s22p4
Na: 1s22s22p63s1
As: 1s22s22p63s23p64s23d104p3

14. ELECTRON CONFIGURATION OF IONS
Cl: [Ne]3s23p5
Cl-: [Ne]3s23p6
Na: [Ne]3s1
Na+: [Ne]
F : 1s22s22p5
F- : 1s22s22p6
Co: [Ar]3d74s2
Co3+: [Ar]3d6

15. EFFECTIVE NECLEAR CHARGE
Negatively charged electrons are attracted to the positively
charged nucleus
The force of attraction between an electron and the nucleus
- Depends on the magnitude of the net nuclear charge acting
on the electrons
- Depends on the average distance between the nucleus and
the electron
(Coulombs law)

16. EFFECTIVE NECLEAR CHARGE
The force of attraction
- Increases with increasing nuclear charge
- Decreases with increasing average distance between electrons
and the nucleus
- Electrons also experience repulsion by other electrons in the atom
Zeff = Z – S
Zeff = effective nuclear charge
Z = actual nuclear charge (number of protons in the nucleus)(> Zeff)
S = screening constant (represents number of core electrons)

17. EFFECTIVE NECLEAR CHARGE
Atomic number of Na = 11
Number of valence electrons = 1
Number of core electrons = 10
As simplified from this model
Z = 11
S = 10
Zeff = = +1
In actual fact
Zeff in Na is about +2.5

18. EFFECTIVE NECLEAR CHARGE
In general
Zeff in s orbital > Zeff in p orbital > Zeff in d orbital > Zeff in f orbital
- This is the result of the trend in energy levels
ns < np < nd < df
- Zeff increases across the periods (from left to right) in the
periodic table (Z increases but S remains the same)
- Zeff changes slightly down the groups in the periodic table
(essentially the same in a given group)

19. SIZES OF ATOMS - Atomic radius tends to decrease across the periods
(from left to right) in the periodic table
- Due to increase in effective nuclear charge which draws
valence electrons closer to the nucleus
- Atomic radius tends to increase down the groups (from top
to bottom) of the periodic table
- Due to increase in principal quantum number of the
outer electrons (number of shells)

20. SIZES OF ATOMS - Cations are smaller than their parent atoms
- Decrease in the number of electrons decreases
electron-electron repulsions
- Anions are larger than their parent atoms
- Increase in the number of electrons increases
- Ionic size increases down the group of the periodic table
for ions carrying the same charge

21. ISOELECTRONIC SERIES
- A group of atoms and ions containing the same number of electrons
Due to the same number of electrons
- Ionic radius decreases with increasing nuclear charge
(electrons are more strongly attracted to the nucleus)
increasing nuclear charge
O2- , F- , Na+ , Mg2+ , Al3+
S2- , Cl- , K+ , Ca2+ , Ga3+
Se2- , Br- , Rb+ , Sr2+ , In3+
decreasing ionic radius

22. IONIZATION ENERGY The energy required to remove an electron from
a gaseous atom or ion
X(g) → X+(g) + e-
E is positive
- The atom or ion is assumed to be in its ground state
- The highest energy electron is always removed first
Units: kJ/mol
kJ/mol = 1 eV

23. IONIZATION ENERGY Ionization Energies of Magnesium (Mg)
Mg(g) → Mg+(g) + e- I1 = 738 kJ/mol
Mg+(g) → Mg2+(g) + e- I2 = 1450 kJ/mol
Mg2+(g) → Mg3+(g) + e- I3 = 7734 kJ/mol
I1 < I2 < I3

24. IONIZATION ENERGY First Ionization Energy (I1)
- Energy required to remove the highest energy electron of an atom
- The first electron is removed from a neutral atom
- The second electron is removed from a positive ion
(more difficult)
- Increase in positive charge binds electrons more tightly
- Large jump is observed in going from removal of valence
electrons to removal of core electrons

25. IONIZATION ENERGY
- First ionization energy increases across the period of the
periodic table (from left to right)
- Electrons added in the same principal quantum number do
not completely shield increasing nuclear charge
- First ionization energy decreases down the group of the
periodic table (from top to down)
- As n increases, the size of orbital increases (distance
from nucleus increases) and electrons are easier to remove

26. IONIZATION ENERGY
- Discontinuities are due to electron repulsions and shielding
(Be to B, N to O)
- Representative elements show a larger range of values
of I1 than the transition-metal elements
- Smaller atoms have higher ionization energies

27. - Several cations with the pseudo-noble-gas electron configuration
PSEUDO-NOBLE-GAS CONFIGURATION
- Several cations with the pseudo-noble-gas electron configuration
are more stable
[noble gas](n-1)d10
In: [Kr]5s24d105p1
In+: [Kr]5s24d10
Pseudo = In3+: [Kr]4d10
Sn: [Kr]5s24d105p2
Sn2+: [Kr]5s24d10
Pseudo = Sn4+: [Kr]4d10

28. ELECTRON AFFINITY - The energy change that occurs when an electron is
added to a gaseous atom
- A measure of the attraction of the atom for the
added electron
- Energy is released when an electron is added to most atoms
X(g) + e- → X-(g)
E is negative
Units: kJ/mol

29. ELECTRON AFFINITY Cl(g) + e- → Cl-(g) E = -349 kJ/mol
- The greater the attraction between an atom and an added
electron the more negative the atom’s electron affinity
- Electron affinities for noble gases are positive values
(E > 0)
- Halogens have the most negative electron affinities
- Electron affinity changes slightly down the group
of the periodic table

30. METALS - Refer to chapter 2 for properties of metals
- Metals tend to have low ionization energies
- Metals form positive ions relatively easily
- Metals lose electrons (oxidize) when they
undergo chemical reactions
Charge on metals
Alkali metals: 1+
Alkaline earth metals: 2+

31. METALS - The charge on transition metals do not follow any
obvious pattern
- Transition metals are able to form more than one positive ion
- Compounds of metals and nonmetals are ionic
2Na(s) + Cl2(g) → 2NaCl(s)
(contains Na+ and Cl- ions)

32. METALS Most Metal Oxides are Basic
- Dissolve in water to form metal hydroxides
Na2O(s) + H2O(l) → 2NaOH(aq)
O2-(aq) + H2O(l) → 2OH-(aq)
(net ionic equation)
- React with acid to form salt and water
NiO(s) + 2HNO3(aq) → Ni(NO3)2(aq) + H2O(l)

33. NONMETALS - Refer to chapter 2 for properties of nonmetals
- Nonmetals tend to have high electron affinities
- Nonmetals tend to gain electrons when they react with metals
- Compounds composed of only nonmetals are generally
molecular substances

34. NONMETALS Most Nonmetal Oxides are Acidic CO2(g) + H2O(l) → H2CO3(aq)
(acidity of rainwater)
- Dissolve in basic solutions to form salt and water
CO2(g) + 2NaOH(aq) → Na2CO3(aq) + H2O(l)

35. ALKALI METALS (GROUP IA)
- Alkali means ‘ashes”
Relatively abundant in the
- earth’s crust (Na, K)
- sea water
- human bodies
- Have low densities and melting points
- Very reactive and readily lose an electron to form 1+ ions

36. ALKALI METALS (GROUP 1A)
- Form hydrides with hydrogen and sulfides with sulfur
2M(s) + H2(g) → 2MH(s)
2M(s) + S(s) → M2S(s)
- React vigorously with water to produce hydrogen gas and
alkali metal hydroxide
(very exothermic and may explode)
2M(s) + 2H2O(l) → 2MOH(aq) + H2(g)

37. ALKALI METALS (GROUP 1A)
- Can react with oxygen to form oxides, peroxides, and superoxides
Li forms only oxides (O2-)
4Li(s) + O2(g) → 2Li2O(s) (lithium oxide)
Na can form peroxides (O22-)
2Na(s) + O2(g) → Na2O2(s) (sodium peroxide)
K, Rb, and Cs can form peroxides (O22-) and superoxides (O2-)
K(s) + O2(g) → KO2(s) (potassium superoxide)

38. ALKALINE EARTH METALS (GROUP 2A)
Compared to alkali metals
- harder
- more dense
- higher melting points
- less reactive than the respective adjacent alkali metal
- Tend to lose two outer s electrons to from 2+ ions
- Give off characteristic colors when heated in a flame
(salts used in fireworks)

39. ALKALINE EARTH METALS (GROUP 2A)
- Reactivity increases down the group
- Beryllium does not react with water
- Magnesium reacts slowly with water
- Calcium and elements below react readily with water
Ca(s) + 2H2O(l) → Ca(OH)2 + H2(g)

40. HYDROGEN - First element in the periodic table
(1s1 electron configuration)
- Nonmetal
- Can be metallic under extreme pressures
- Colorless diatomic gas
- Has very high ionizaton energy
- More than double those of alkali metals
- Due to absence of nuclear shielding of the 1s electron

41. HYDROGEN - Does not easily lose its valence electron
- Share with nonmetals to form molecular compounds
- Can lose its electron to form a cation (H+)
- Can gain electron to form the hydride ion (H-)

42. CHALCOGENS (GROUP 6A) THE OXYGEN GROUP
- Properties change from nonmetallic to metallic down the group
Nonmetallic properties: oxygen, sulfur, selenium
Metallic properties: tellurium and below
- Oxygen is a colorless gas at room temperature
- The other group members are solids
- Oxygen exists as O2 (oxygen gas) and O3 (ozone)
- Allotropes (different forms of the same element in the same state)

43. CHALCOGENS (GROUP 6A) THE OXYGEN GROUP
- O2 can produce O3 in lightning storms
3O2(g) → 2O3(g) Ho = kJ
- Sulfur also has several allotropic forms
- The most common is S8 (yellow solid)

44. THE HALOGENS (GROUP 7A) - All halogens are nonmetals
- Melting and boiling points increase with increasing
atomic number
- Consist of diatomic molecules
(F2, Cl2, Br2, and I2)
- Form colored gases

45. THE HALOGENS (GROUP 7A) At room temperature
- Fluorine and chlorine are gases
- Bromine is a liquid
- Iodine is a solid
- Have highly negative electron affinities
- Tend to gain electrons to form 1- ions
- Reactivity decreases down the group

46. THE HALOGENS (GROUP 7A)
- React readily with most metals to form ionic halides
- React with hydrogen to form gaseous hydrogen halides
H2(g) + X2 → 2HX(g)
- Hydrogen halides dissolve in water to form acids [HCl(aq)]
- Fluorine is very reactive (dangerous)
- Chlorine is the most industrially useful

47. THE NOBLE GASES (GROUP 8A)
- Nonmetals
- Monatomic
- Gases at room temperature
- Have completely filled s and p subshells
- Have high first ionization energies
- Have stable electron configuration
- Unreactive