Periodic Properties of the Elements

Contributed by:

Jonathan James Fri, Jan 14, 2022 07:13 PM UTC

The highlights are:
1. Development of Periodic Table
2. Effective nuclear charge
3. Periodic trends
4. Properties of Metals, Non-Metals, and Metalloids

1. Chapter 7
Periodic Properties
of the Elements
Periodic
Properties
of the
Elements

2. Development of Periodic Table
• Elements in the
same group
generally have
similar chemical
properties.
• Properties are not
identical, however.
Periodic
Properties
of the
Elements

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

4. Development of Periodic Table
Mendeleev, for instance, in 1871 predicted germanium
(which he called eka-silicon) to have an atomic weight Periodic
Properties
between that of zinc and arsenic, but with chemical of the
properties similar to those of silicon. Elements

5. Development of Periodic Table
Mendeleev’s prediction was on the money
Periodic
Properties
But why? (Mendeleev had no clue). of the
Elements

6. Periodic Trends
• In this chapter we’ll explain why
• We’ll then rationalize observed trends in
Sizes of atoms and ions.
Ionization energy.
Electron affinity.
Periodic
Properties
of the
Elements

7. Effective Nuclear Charge
• In a many-electron atom,
electrons are both attracted to
Na atom looks like this: the nucleus and repelled by
other electrons.
• The nuclear charge that an
electron “feels” depends on
both factors.
• It’s called Effective nuclear
charge.
• electrons in lower energy
levels “shield” outer electrons
from positive charge of
nucleus.
Periodic
Properties
of the
Elements

8. Effective Nuclear Charge
The effective nuclear
charge, Zeff, is:
Zeff = Z − S
Where:
Z = atomic number
S = screening constant,
usually close to the
number of inner (n-1)Periodic
electrons. Properties
of the
Elements

9. Effective Nuclear Charge
• Example: Which element’s outer shell or
“valence” electrons is predicted to have the
largest Effective nuclear charge? Kr, Cl or O?
Periodic
Properties
of the
Elements

10. Effective Nuclear Charge
• Example: Which element’s outer shell or
“valence” electrons is predicted to have the
largest Effective nuclear charge? Kr, Cl or O?
• Cl: Zeff ≈ 17 - 10 = 7
• O: Zeff ≈ 8 - 2 = 6
• N: Zeff ≈ 7 - 2 = 5
• Ca: Zeff ≈ 20 - 18 = 2
Periodic
Properties
of the
Elements

11. Valence electrons
Many chemical properties depend on the valence electrons.
Valence electrons: The outer electrons, that are involved in
bonding and most other chemical changes of elements.
Rules for defining valence electrons.
1. In outer most energy level (or levels)
2. For main group (representative) elements (elements in s
world or p world) electrons in filled d or f shells are not
valence electrons
3. For transition metals, electrons in full f shells are not
valence electrons.
Periodic
Properties
of the
Elements

12. Valence electrons
Many chemical properties depend on the valence electrons.
Valence electrons: The outer electrons, that are involved in bonding and most
other chemical changes of elements.
Rules for defining valence electrons.
1. In outer most energy level (or levels)
2. For main group (representative) elements (elements in s world or p world)
electrons in filled d or f shells are not valence electrons
3. For transition metals, electrons in full f shells are not valence electrons.
Examples: (valence electrons in blue)
P: [Ne]3s23p3
As: [Ar] 4s23d104p3
I: [Kr]5s24d105p5
Ta: [Kr]6s24f145d3
Zn: [Ar]4s23d10 Periodic
Properties
of the
Elements

13. Sizes of Atoms
The bonding atomic
radius is defined as
one-half of the
distance between
covalently bonded
Periodic
Properties
of the
Elements

14. Sizes of Atoms
inc n
rea n g
sin i
gE eas
ff
ncr
i
Bonding atomic radius tends to… Periodic
Properties
…decrease from left to right across a row due to increasing Zeff. of the
Elements
…increase from top to bottom of a column due to increasing value of n

15. Exams are being graded now but aren’t finished.
We may have them by the end of today.
My T.A.s had to sacrifice bunnys for science.
Periodic
Properties
of the
Elements

16. Sizes of Ions
Ionic size
depends upon:
Nuclear charge.
Number of
electrons.
Orbitals in which
electrons reside.
Periodic
Properties
of the
Elements

17. Sizes of Ions
• Cations are
smaller than their
parent atoms.
 The outermost
electron is
removed and
repulsions are
reduced.
Periodic
Properties
of the
Elements

18. Sizes of Ions
• Anions are larger
than their parent
atoms.
 Electrons are
added and
repulsions are
increased.
Periodic
Properties
of the
Elements

19. Sizes of Ions
• Ions increase in size
as you go down a
column.
 Due to increasing
value of n.
Periodic
Properties
of the
Elements

20. Sizes of Ions
• In an isoelectronic series, ions have the same
number of electrons.
• Ionic size decreases with an increasing
nuclear charge.
Periodic
Properties
of the
Elements

21. atom/ion size examples
• Put the following in order of size, smallest to
largest:
• Na, Na+, Mg, Mg2+, Al, Al3+, S, S2-, Cl, Cl-
Periodic
Properties
of the
Elements

22. Atom size examples
Al3+, Mg2+, Na+, Cl, S, Al, Mg, Na, Cl-, S2-
Start with atoms with no n=3 electrons, order isoelectronic by nuclear charge.
Next, neutral atoms highest Eff first
Last, anions, highest Eff first
Ambiguity: anions versus neutrals (is Cl- really larger than Na?)
Don’t worry about it.
Periodic
Properties
of the
Elements

23. Ionization Energy
• Amount of energy required to remove
an electron from the ground state of a
gaseous atom or ion.
First ionization energy is that energy
required to remove first electron.
Second ionization energy is that energy
required to remove second electron, etc.
El -------> El+ + e-
Na -------> Na+ + e-
Periodic
Properties
of the
Elements

24. 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.
Periodic
Properties
of the
Elements

25. Trends in First Ionization Energies
• going down a
column, less
energy 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.
Periodic
Properties
of the
Elements

26. Trends in First Ionization Energies
• Generally, it gets
harder to remove an
electron going
across.
 As you go from left to
to right, Zeff
increases.
Periodic
Properties
of the
Elements

27. Trends in First Ionization Energies
On a smaller
scale, there
are two jags
in each line.
Why?
Periodic
Properties
of the
Elements

28. Trends in First Ionization Energies
• The first occurs
between Groups IIA
and IIIA.
• Electron removed from
p-orbital rather than s-
orbital
 Electron farther from
nucleus
 Small amount of
repulsion by s Periodic
electrons. Properties
of the
Elements

29. Trends in First Ionization Energies
• The second occurs
between Groups VA
and VIA.
 Electron removed
comes from doubly
occupied orbital.
 Repulsion from other
electron in orbital
helps in its removal.
versus:
Periodic
Properties
of the
Elements

30. Electron Affinity
Energy change accompanying addition of
electron to gaseous atom:
Cl + e−  Cl−
Periodic
Properties
of the
Elements

31. Trends in Electron Affinity
In general, electron affinity becomes more Periodic
exothermic as you go from left to right across Properties
of the
a row. Elements

32. Trends in Electron Affinity
There are also two
discontinuities in this
trend.
Periodic
Properties
of the
Elements

33. Trends in Electron Affinity
• The first occurs
between Groups IA
and IIA.
 Added electron must
go in p-orbital, not s-
orbital.
 Electron is farther
from nucleus and
feels repulsion from
s-electrons. Periodic
Properties
of the
Elements

34. Trends in Electron Affinity
• The second occurs
between Groups IVA
and VA.
 Group VA has no
empty orbitals.
 Extra electron must
go into occupied
orbital, creating
repulsion.
Periodic
Properties
of the
Elements

35. Properties of Metals, Nonmetals,
and Metalloids
Periodic
Properties
of the
Elements

36. Metals versus Nonmetals
Differences between metals and nonmetals
tend to revolve around these properties.
Periodic
Properties
of the
Elements

37. Metals versus Nonmetals
• Metals tend to form cations.
• Nonmetals tend to form anions.
The common elemental ions
Note ions in s and p world all result from filling or empyting
a subshell. Periodic
Properties
What about the transition metals? What’s going on there? of the
Elements

38. Transition Metal ions
Note: many have +2 charge.
They actually lose all their ns electrons first!
Mn --> Mn2+: [Ar]4s23d5 ---> [Ar]3d5
Cu --> Cu+ [Ar]4s23d9 ---> [Ar]3d10
Periodic
Properties
of the
Elements

39. Tend to be lustrous,
malleable, ductile,
and good
conductors of heat
and electricity.
Periodic
Properties
of the
Elements

40. Metals
• Compounds formed
between metals and
nonmetals tend to
be ionic.
• Metal oxides tend to
be basic.
Periodic
Properties
of the
Elements

41. • Dull, brittle
substances that are
poor conductors of
heat and electricity.
• Tend to gain
electrons in
reactions with
metals to acquire
noble gas
configuration. Periodic
Properties
of the
Elements

42. Nonmetals
• Substances
containing only
nonmetals are
molecular
compounds.
• Most nonmetal
oxides are acidic.
Periodic
Properties
of the
Elements

43. • Have some
characteristics of
metals, some of
nonmetals.
• For instance, silicon
looks shiny, but is
brittle and fairly poor
conductor.
Periodic
Properties
of the
Elements

44. Group Trends
Periodic
Properties
of the
Elements

45. Alkali Metals
• Soft, metallic solids.
• Name comes from
Arabic word for
ashes.
Periodic
Properties
of the
Elements

46. Alkali Metals
• Found only as compounds in nature.
• Have low densities and melting points.
• Also have low ionization energies.
Periodic
Properties
of the
Elements

47. Alkali Metals
Their reactions with water are famously exothermic.
Periodic
Properties
of the
Elements

48. Alkali Metals
• Alkali metals (except Li) react with oxygen to
form peroxides.
• K, Rb, and Cs also form superoxides:
K + O2  KO2
• Produce bright colors when placed in flame.
Periodic
Properties
of the
Elements

49. Alkaline Earth Metals
• Have higher densities and melting points than
alkali metals.
• Have low ionization energies, but not as low
as alkali metals. Periodic
Properties
of the
Elements

50. Alkaline Earth Metals
• Be does not react
with water, Mg
reacts only with
steam, but others
react readily with
water.
• Reactivity tends to
increase as go down
group. Periodic
Properties
of the
Elements

51. Group 6A
• Oxygen, sulfur, and selenium are nonmetals.
• Tellurium is a metalloid.
• The radioactive polonium is a metal.
Periodic
Properties
of the
Elements

52. • Two allotropes:
 O2
 O3, ozone
• Three anions:
 O2−, oxide
 O22−, peroxide
 O21−, superoxide
• Tends to take electrons
from other elements
(oxidation) Periodic
Properties
of the
Elements

53. Sulfur
• Weaker oxidizing
agent than oxygen.
• Most stable
allotrope is S8, a
ringed molecule.
Periodic
Properties
of the
Elements

54. Group VIIA: Halogens
• Prototypical nonmetals
• Name comes from the Greek halos and
gennao: “salt formers”
Periodic
Properties
of the
Elements

55. Group VIIA: Halogens
• Large, negative electron
affinities
 Therefore, tend to oxidize
other elements easily
• React directly with metals
to form metal halides
• Chlorine added to water
supplies to serve as
disinfectant Periodic
Properties
of the
Elements

56. Group VIIIA: Noble Gases
• Astronomical ionization energies
• Positive electron affinities
 Therefore, relatively unreactive
• Monatomic gases Periodic
Properties
of the
Elements

57. Group VIIIA: Noble Gases
• Xe forms three
compounds:
 XeF2
 XeF4 (at right)
 XeF6
• Kr forms only one stable
compound:
 KrF2
• The unstable HArF was Periodic
Properties
synthesized in 2000. of the
Elements

58. Exam 2 review:
• Chapter 5, thermochemistry
• Chapter 6, atomic structure
• Chapter 7, periodic trends
Periodic
Properties
of the
Elements

59. Exam 2 review:
• Chapter 5
 Heat vs. work, nature of energy
• potential energy vs. kinetic energy
• nature of temperature
• system versus surroundings
 1st law of thermodynamics:
 E calculations
• example: Calculate the change in internal energy and whether the process is endo or exo thermic:
– 100g of water is cooled from 90 °C to 40 °C.
 Enthalpy of reaction. Using stoichiometry and enthalpy of reaction to calculate things:
• example problem:
• Ag+(aq) + Cl-(aq) -------> AgCl(s) H = -65.5 kJ
a. Calculate H for the formation of 2.5 g of AgCl
Periodic
Properties
of the
Elements

60. •
Exam 2 review:
Ag+(aq) + Cl-(aq) -------> AgCl(s) H = -65.5 kJ
Calculate H for the formation of 2.5 g of AgCl
MW. 143.319 g/mol (AgCl).
2.5 g/143.319 g/mol = 0.0174 mole
0.0174 mol(-65.5 kJ/mol AgCl) = -1.14 kJ
 Calorimetry problem:
• example problem:
10 g of NaOH is dissolved in 100 mL of water in a calorimeter, the temperature changes from 23.6 to 47.4 °C.
Calculate H for the process, assume the specific heat of the solution is 4.184 j/°Kg, the same as water.
Periodic
Properties
of the
Elements

61. Chapter 5
q = specific heat(g solution)(T)
q = 4.184j/Kmol(110 g)(26.4 -43.3) = -10953 J
moles NaOH = 10g/40 g = .25 mole
H = -10953 J/.25 mol = -43812 = 40000 J/mol NaOH.
 Hess’s law:
• Given a series of reactions, rearrange to find H for the reaction in question:
• Example problem:
• Given the data:
• N2(g) + O2(g) ----> 2NO(g) H =180.7 kJ
• 2NO(g) + O2(g) -----> 2NO2 H = -113.1 kJ
• 2N2O(g) -----> 2N2(g) + O2(g) H = -163.2 kJ
• Calculate: N2O(g) +NO2(g) -----> O2(g)
 Enthalpies of formation (the tables of enthalpies of formation):
 Hprdts - Hreactants = Hreaction
Periodic
Properties
of the
Elements

62. Chapter 5
 Hess’s law:
• Given a series of reactions, rearrange to find H for the reaction in question:
• Example problem:
• Given the data:
• N2(g) + O2(g) ----> 2NO(g) H =180.7 kJ
• 2NO(g) + O2(g) -----> 2NO2 H = -113.1 kJ
• 2N2O(g) -----> 2N2(g) + O2(g) H = -163.2 kJ
• Calculate: N2O(g) +NO2(g) -----> 3NO(g)
• NO2 ---- NO(g) + 1/2O2(g) 113.1/2
• N2O(g) -----> N2(g) + 1/2O2(g) H = -163.2/2 kJ
• N2(g) + O2(g) ----> 2NO(g) H =180.7 kJ
• N2O(g) +NO2(g) -----> 3NO(g) H =155.6 kJ
Periodic
Properties
of the
Elements

63. Chapter 5
 Enthalpies of formation (the tables of enthalpies of formation):
 Hprdts - Hreactants = Hreaction
Periodic
Properties
of the
Elements

64. Chapter 6
• Chapter 6
 Characteristics of waves (v = ):
• What is wavelength?
• What is frequency?
 Electromagnetic radiation:
• E = h
• visible spectrum (ROYGBV)
• rest of spectrum
Periodic
Properties
of the
Elements

65. Chapter 6
• Chapter 6
 Characteristics of waves (v = ):
 Black body radiation
 Photo-electric effect
 Heisenberg uncertainty: (mvx ≥ h)
 Line spectra of atoms
 matter waves (De Brogli)
v = λν
v
ν=
λ
2 v
E = mv = hν = h
λ
h
λ= Periodic
mv Properties
of the
Elements

66. Chapter 6
• Chapter 6
 Wavefunctions and quantum mechanics
• wavefunction vs. probability distribution
• orbitals and quantum numbers
 Quantum numbers
• what are the four?
– principle (energy) n = 1,2,3...
– azimuthal (shape) l = 0,1, 2... n-1
– magnetic (orientation) ml = -l,...0,...+l
– spin (differentiates two electrons in same orbital) (±1/2)
• naming the l qm:
• l=0, s, l=1, p, l=2, d, l=3, f
 Shapes of orbitals
Periodic
Properties
of the
Elements

67. Chapter 6
• Chapter 6
 Many electron atoms
 Energy of orbitals in H versus other atoms with other electrons.
Periodic
Properties
of the
Elements

68. Chapter 6
• Chapter 6
 Pauli exclusion principle
 Hund’s rule (don’t pair until you have to)
 Electron configurations
Periodic
Properties
of the
Elements

69. Chapter 7
• Periodic trends
• Effective nuclear charge
• trends in atomic radius
• trends in ion radius
• Ionization energy, trends
• electron affinity, trends.
Periodic
Properties
of the
Elements

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