1. The Periodic Table and Periodic Law
Chemistry I Chapter 6
The Periodic Table and Periodic Law

2. 1.)Who was Antione Laviosier?
Known as the Father of Modern Chemistry
In the 1790’s he created a list of the 23 known chemicals and their properties.

3. 2.) Who John Newlands?
An English Chemist ( ) who noticed a that when the elements are arranged in order of increasing atomic mass the properties repeated every 8th element.
He called this relationship the Law of Octaves

4. 3.What are two reasons that Newland’s Law of Octaves was not accepted?
1st – the law of octaves did not work for all elements.
2nd by comparing chemicals to musical notes many fellow scientists believed that Newlands was being unscientific.

5. 4.In 1869 which two scientists demonstrated a connection between atomic mass and elemental properties. Who were these two scientists?
Dimtri Mendeleev and Lothar Meyer

6. 5.How did Mendeleev arrange the periodic table?
In order of increasing atomic mass.

7. 6. Mendeleev’s periodic table still had problems
6.Mendeleev’s periodic table still had problems. When the elements were placed in increasing atomic masses, several of the elements did not “fit” correctly. How was this issued resolved in 1913?

8. 6.)In 1913 English Chemist Henry Moseley discovered that each element had a unique number of protons in the nucleus. This the atomic number. When all of the elements are placed in order of increasing atomic number all of the inconsistencies were resolved.

9. 7.What is the periodic law?
There is a periodic repetition of chemical and physical properties of elements when they are placed in order of increasing atomic number.

10. The Modern Periodic Table

11. 8.In the modern periodic table, how are the elements arranged?
By increasing atomic number

12. 9.What do we call the columns of elements?
Groups or families

13. 10.What do we call the horizontal rows of elements?
periods

14. 11.The groups designated with an “A” are called what?
The representative elements
1A – 8A
Also called the main group elements because they represent a wide range of chemical and physical properties.

15. 12.What are the elements designated with a “B” referred to as?
As the transition elements.
The two groups at the bottom of the period table are known as the inner transition elements. The top group are the lanthanides, the bottom group are the actinides.

16. The scientific community has recently decided to number the groups However there may be situations in which you are expected to know the A and B designations.

17. 13.What are the three classifications of elements?
Metals, nonmetals and metalloids ( also known as semi-metals)

18. 14.What are the properties of metals?
Good conductors of heat and electricity
High luster ( shiny)
Malleable
Ductile
High melting points

19. 15.What are the properties of non-metals?
Poor conductors of heat and electricity ( they are insulators)
Brittle
Low melting points
Dull

20. 16.Where are the metalloids? Be able to name the metalloids.
The metalloids are found on either side of the stair step line separating the metals and nonmetals.
The metalloids include:B, Si,Ge,As,Sb,Te,Po,and At ( The only element that borders this line that is not a metalloid is Aluminum)

21. 17.Name each of the following groups
Group 1 ( 1A) Alkali metals
Group 2 (2A) Alkaline Earth metals
Group 13 (3A) Boron Family
Group 14 ( 4A) Carbon Family
Group 15 (5A) Nitrogen Family
Group 16 (6A) Oxygen Family
Group 17 (7A) Halogens
Group 18 (8A) Noble Gases

22. 18.What are the Group B elements called?
The transition elements (metals)

23. Section 2: Classification of the Elements

24. Last chapter we learned how to write out electron configurations and noble gas configurations. The electron configuration determines the chemical properties of the element. By using the periodic table you can determine both the electron configuration and the number of valence electrons.

25. 19.How many valence electrons does each group have?
Group (with the exception of He which has 2)

26. 20.What is the group configuration of each family of the representative elements ( 1A to 8A)
Group 1: ns Group 2: ns 2
Group 13: ns2np1 Group 14: ns2np2
Group 15: ns2np3 Group 16:ns2np4
Group 17: ns2np Group 18:ns2np6

27. 21.What is an electron dot structure ( also known as Lewis dot diagrams)
A method used to show valence electrons. The chemical symbol is used to represent all of the filled inner electron orbitals. The number of valence electrons is shown by using dots. Only two dots per side. Each side gets one dot before any get two ( with the exception of He)

28. 22.)Draw the Electron dot structure for:
Li Mg
Al Si
N O
Cl Ar

29. 23.What are the four blocks on the periodic table?
S block – alkali metals (1A) and Alkaline Earth metals (2A)
s orbitals are filling
P block – groups
p orbitals are filling

30. 23 cont’d) What are the four blocks on the periodic table?
d block – transition metals
d orbitals are filling
f block – Lanthanides and actinides ( inner transition metals)
f orbitals are filling

31. Practice problems page 181 and 186

32. Periodic Trends
What are the periodic trends that can be predicted with the use of the periodic table?

33. 24.What are the periodic trends that can be predicted with the use of the periodic table?
Atomic radius
Ionic size
Ionization energy
Electron affinity
Electronegativity

34. 25.What is the trend for atomic radius?
As you go down the periodic table, the atomic radius increases
As you go across the periodic table from left to right the atomic radius decreases.

35. 26.Why would atomic radius decrease as you move from left to right across the period?
The attraction between the electron cloud and the nucleus grows because we add more protons and electrons without increasing the distance. The increased electrical attraction shrinks the electron cloud.

36. Practice problems pg 189

37. 27.What is an ion?
An atom that has lost or gained an electron. When atoms lose electrons they become positively charged; when they gain an electron they become negatively charged.

38. 28.What is the trend for ionic radius?
When atoms gain electrons the ion becomes larger than the atom.
When atoms lose electrons the ion becomes smaller than the atom.
Continued ( 1 more slide)

39. 28.What is the trend for ionic radius?
Overall, when you move left to right across the periodic table, the ionic size decreases.
( exception, when you transition from positive ions to negative ions at the nitrogen family, ionic size increases.)
When you move down the periodic table ionic size increases.

40. 29.Why are negatively charged ions much larger than the parent atom?
There are more electrons to ‘hold on’ to than there are protons. Each electron is being held with less electrical attraction. In addition the electrons repel each other.

41. 30.Why are positively charged ions much smaller than the parent atom?
There are more protons than electrons. Each electron has a greater electrical attraction acting on it. In addition because there are fewer electrons there is less repulsive force between electrons.

42. 31.What is ionization energy?
The energy needed to remove an electron from an atom. Even though an atom may be more stable if it loses an electron and becomes an ion, it still takes energy to remove the electron from the atom.

43. 32.What is the difference between a 1st ionization energy value and successive ionization values?
The 1st ionization value is the energy needed to remove just the 1st electron. Successive ionization energies are the energies needed to remove the 2nd, 3rd, 4th etc. electron.

44. 33.What does a high ionization energy indicate?
The atom has a very strong hold on the electron.

45. 34.What does a low ionization energy indicate?
The electrons are being held fairly loosely.

46. Look at figure 6-16 on page 191. What do you notice about the highest first ionization energy for each element in the group?
They are all noble gases. Noble gases do not want to lose electrons.

47. 35.What do you notice about the lowest ionization energy for each group?
They are all alkali metals (1A) It takes a little energy to remove an electron from an alkali metal, but relatively little energy because the alkali metals will be more stable as an ion with a plus one charge.

48. 36.How does the first ionization value for hydrogen compare to rubidium?
Rb is much lower than hydrogen

49. 37.What is the general trend for first ionization energies?
It decreases as you go down the periodic table. ( atoms get larger so it is easier to remove an electron)
It increases as you go across the periodic table left to right. (the atoms get smaller so it is more difficult to remove an electron)

50. 38.What can successive ionization energies tell us?
Look at table 6-2 page 192.
The ionization energy at which the large jump occurs is related to the number of valence electrons.

51. 38.) What can successive ionization energies tell us?
Li has 1 valence electron- the big jump happens after the 1st ionization.
Be has 2 v.e. – the big jump happens after the 2nd ionization.
B has 3 v.e. – the big jump happens after the 3rd ionization.

52. The ion of an atom of a representative element will have the same electron configuration as the noble gas before it..

53. 39.Write the electron configuration for Na and the Na+
Na - 1s22s22p63s1
Na+ - 1s22s22p6
Ne - 1s22s22p6

54. Write the electron configuration for Cl and Cl-
Cl - 1s22s22p63s23p5
Cl- - 1s22s22p63s23p6
Ar - 1s22s22p63s23p6

55. 40. What is the octet rule?
Atoms tend to gain, lose or share electrons in order to acquire a full valence shell which is 8 electrons( with the exception of H and He which only need 2 valence electrons to have a full shell)

56. The octet rule is useful for determining the type of ion that will form and how electrons will be shared in covalent bonds.

57. 41.What is electronegativity?
It is a measure that indicates the ability of an atom to attract the electrons in a chemical bond.

58. 42.How are electronegativity values expressed?
In units called Paulings that have been calculated. The units range from 0 to 4 with 4 being the largest electronegativity value.

59. 43.What is the general trend in electronegativity values.
As you go down the periodic table the electronegativity values decrease. As you move across the periodic table to the halogens the electronegativity values increase. (the exception is the noble gases which do not want to gain or lose electrons.)

60. Electronegativity differences are mainly used to determine if a bond is covalent ( electrons are shared) or is ionic ( one atom takes one or more electrons away from another) If two atoms have an electronegativity difference of more than 1.8 the bond is considered ionic.

61. Using the electronegativity chart, determine if the following bonds are ionic or covalent.

62. A difference of greater than 1. 8 makes the bond ionic
A difference of greater than 1.8 makes the bond ionic. One atom is “strong” enough to take the electron away from the other atom. If they are about equal in electronegativity ( less than 1.8) then they share electrons (covalent bond)

63. Na-Br
C-O
O-O
N-H
Na -Cl