Electronic Structure of Atoms & Periodic Table

1. CHEMISTRY - DMCU 1233 Fakulti Kejuruteraan Mekanikal, UTeM Lecturer: IMRAN SYAKIR BIN MOHAMAD MOHD HAIZAL BIN MOHD HUSIN NONA MERRY MERPATI MITAN Electronic Structure of Atoms & Periodic Table Chapter 4

2. 4.1 HISTORY OF ATOMIC MODEL

3. 4.1 HISTORY OF ATOMIC MODEL Plum Pudding Model Billiard Ball Model Solar System Model Bohr Model

4. 4.2 Quantum Numbers A quantum number describes the energies of electrons in atoms • The Bohr model was 1-D model that used one quantum number to describe the distribution of electrons in the atom that representative of the size of the orbit, which was described by the principal quantum number (n). • Meanwhile Schrödinger's model allowed the electron to occupy in 3-D space to describe the orbitals in which electrons can be found. • Each electron in an atom is described by four different quantum numbers. • The first three quantum number from Schrödinger's wave equations are the principal (n), angular (l), and magnetic (ml) quantum numbers describe the size, shape, and orientation in space of the orbitals on an atom. • The fourth quantum number spin (ms) specifies how many electrons can occupy that orbital.

5. 4.2 Quantum Numbers Principal quantum number – ( n ) Angular momentum quantum number – ( l) Magnetic quantum number – ( ml) Spin quantum number – ( ms)

6. n=3 n =1 n=2 Quantum Numbers (n, l, ml, ms) Principal quantum numbern n= 1, 2, 3, 4, …. • Specifies the energy of an electron and the size of the orbital (the distance from the nucleus of the peak in a radial). • All orbitals that have the same value of n are said to be in the same shell (level) • The total number of orbitals for a given n value is n2

7. Quantum Numbers (n, l, ml, ms) Angular momentum quantum numberl for a given value of n,l= 0, 1, 2, 3, … n-1 • Specifies the shape of an orbital with a particular principal quantum number. • The secondary quantum number divides the shells into smaller groups of orbitals called subshells (sublevels). l = 0 s orbital l = 1 p orbital l = 2 d orbital l = 3 f orbital n = 1, l = 0 n = 2, l = 0 or 1 n = 3, l = 0, 1, or 2 Shape of the “volume” of space that the e- occupies

8. l= 0 (s orbitals) l= 1 (p orbitals)

9. l= 2 (d orbitals)

10. l= 3 (f orbitals)

11. Quantum Numbers (n, l, ml, ms) Magnetic quantum numberml for a given value of lml= -l, …., 0, …. +l • Specifies the orientation in space of an orbital of a given energy (n) and shape (l). • This number divides the subshell into individual orbitals which hold the electrons. • There are 2l+1 orbitals in each subshell. for l = 0 (s orbital)ml= 0 if l = 1 (p orbital),ml= -1, 0, or +1 if l = 2 (d orbital),ml= -2, -1, 0, +1, or +2 orientation of the orbital in space

12. ml = -1 ml = 0 ml = 1 ml = -2 ml = -1 ml = 0 ml = 1 ml = 2

13. ORIENTATION OF THE ORBITAL IN SPACE

14. Quantum Numbers (n, l, ml, ms) Spin quantum numberms ms= +½or -½ • Specifies the orientation of the spin axis of an electron. • An electron can spin in only one of two directions (sometimes called up and down). ms = +½ ms = -½ Experimental arrangement for demo the spinning motion of electrons Q & A

16. Relation between quantum number, atomic orbital and number of an electron

17. Quantum Numbers (n, l, ml, ms) Existence (and energy) of electron in atom is described by its unique Quantum Numbers Pauli exclusion principle (Wolfgang Pauli, Nobel Prize 1945) No two electrons in the same atom can have identical values for all four of their quantum numbers • Two electrons in the same orbital must have opposite spins. • Because an electron spins, it creates a magnetic field, which can be oriented in one of two directions. • For two electrons in the same orbital, the spins must be opposite to each other; the spins are said to be paired

18. 2p 2p • The substances are not attracted to magnets and are said to be diamagnetic. • Atoms with more electrons that spin in one direction than another contain unpaired electrons. These substances are weakly attracted to magnets and are said to be paramagnetic. Diamagnetic Paramagnetic all electrons paired unpaired electrons

19. How many electrons can an orbital hold? Quantum Numbers (n, l, ml, ms) Shell – electrons with the same value of n Subshell – electrons with the same values of nandl Orbital – electrons with the same values of n, l, andml

20. How many electrons can be placed in the 3d subshell? How many 2p orbitals are there in an atom? Q & A

21. number of electrons in the orbital or subshell principal quantum number n angular momentum quantum number l 1s1 4.3 Electron configuration Electron configuration is how the electrons are distributed among the various atomic orbitals in an atom. 1s1 Orbital diagram is shows the spin of the electron H

22. Order of orbitals (filling) in multi-electron atom “Fill up” electrons in lowest energy orbitals Aufbau principle - electrons fill orbitals starting at the lowest available (possible) energy states before filling higher states. 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f

23. Outermost subshell being filled with electrons The order in which the electrons are filled in can be read from the periodic table in the following fashion

24. C 6 electrons C 1s22s22p2 B 5 electrons B 1s22s22p1 Li 3 electrons Li 1s22s1 H 1 electron H 1s1

25. The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins (Hund’s rule). Ne 10 electrons Ne 1s22s22p6 F 9 electrons F 1s22s22p5 O 8 electrons O 1s22s22p4 N 7 electrons N 1s22s22p3

26. What is the electron configuration of Mg? What are the possible quantum numbers for the last (outermost) electron in Cl?

27. Electron Configurations of Cations and Anions Of Representative Elements Na [Ne]3s1 Na+ [Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration. Ca [Ar]4s2 Ca2+ [Ar] Al [Ne]3s23p1 Al3+ [Ne] H 1s1 H- 1s2 or [He] Atoms gain electrons so that anion has a noble-gas outer electron configuration. F- 1s22s22p6 or [Ne] F 1s22s22p5 O2- 1s22s22p6 or [Ne] O 1s22s22p4 N3- 1s22s22p6 or [Ne] N 1s22s22p3

28. -1 -2 -3 +1 +2 +3 Cations and Anions Of Representative Elements

29. What neutral atom is isoelectronic with H- ? F-: 1s22s22p6 or [Ne] Na+: [Ne] Al3+: [Ne] O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne] Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne

30. Electron Configurations of Cations of Transition Metals When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals. Fe: [Ar]4s23d6 Mn: [Ar]4s23d5 Fe2+: [Ar]4s03d6 or [Ar]3d6 Mn2+: [Ar]4s03d5 or [Ar]3d5 Fe3+: [Ar]4s03d5 or [Ar]3d5

31. Ion charges

32. Q & A session Name the orbital described by the following quantum numbers : • n = 3, l = 0 • n = 3, l = 1 • n = 3, l = 2 • n = 5, l = 0

33. Q & A session Give the n and l values for the following orbital     a. 1s    b. 3s    c. 2p    d. 4d    e. 5f What and the possible ml values for the following types of orbital?     a. s    b. p    c. dd. f

34. Q & A session How many possible orbital are there for n =     a. 4    b. 10 How many electrons can inhabit all of the n = 4 orbital? Place the following orbital in order of increasing energy:     1s, 3s, 4s, 6s, 3d, 4f, 3p, 7s, 5d, 5p

35. Q & A session Write electron configurations for the following atoms:     a. H     b. Li+     c. N     d. F-     e. Ca

36. Q & A session Draw an orbital diagrams for atoms with the following electron configurations: 1s22s22p63s23p3

38. When the Elements Were Discovered

39. 4.4 HISTORY OF THE PERIODIC TABLE • Antoine Lavoisier (1743–1794) • Classify elements into four groups including light and heat, into metals and non-metals.

40. 4.4 HISTORY OF THE PERIODIC TABLE • Johann Dobereiner (1780–1849) • The first significant groupings of elements by place certain elements in groups of three known as The Law of Triads. • Founded that strontium had about the average properties of calcium and barium, and grouped these three together. • Several more triad groups, including the halogen triad of chlorine, bromine, and iodine, and the alkali metal triad of lithium, sodium, and potassium. • However, due to the inaccuracy of many measurements, including atomic weight, the relationship between large element groups could not be exacted

41. 4.4 HISTORY OF THE PERIODIC TABLE • John Newlands (1837–1898) • arranged known elements horizontally in the ascending order of their atomic masses • Each row consisted of seven elements. • Founded that elements with similar properties • repeated at every eighth element. • This arrangement was known as the Law of Octaves • However, this law was only obeyed by the first 17 elements. • There were no positions allocated for elements yet to be discovered

42. 4.4 HISTORY OF THE PERIODIC TABLE • Lothar Meyer (1830–1895) • Plotted a graph of atomic volume against atomic mass for all known elements. • Founded that elements with the same chemical properties occupied the same relative positions on the curve. • Showed that the properties of the elements were in a periodic pattern with their atomic masses. • Proved that the properties of the elements recur periodically.

43. 4.4 HISTORY OF THE PERIODIC TABLE • Dmitri Mendeleev (1834–1907) • Showed that the properties of elements changed periodically with their atomic mass. • Arranged the elements in the order of increasing atomic mass and grouped them according to similar chemical properties. • Able to predict the properties of undiscovered elements and left gap for these elements • For examples correctly predicted the properties of the elements gallium, scandium and germanium which were only discovered later • Mendeleevs table was used as a blueprint for the modern periodic table Mendeleev’s periodic table

44. 4.4 HISTORY OF THE PERIODIC TABLE • Henry J. G. Moseley (1887–1915) • Based on the x-ray spectrum of elements studies, he concluded that the proton numbers should be used as a basis for the periodic change of chemical properties instead of the atomic mass. • Rearranged the elements in the ascending order of their proton numbers • Similar to Mendeleev, Moseley left gaps for elements yet to be discovered. • produced a periodic table which was almost the same as Mendeleevs periodic table. • Due to Moseley’s work, the periodic table was successfully developed and being used today. • The modern periodic table is based on the arrangement of elements in the ascending order of their proton numbers.

45. 4.4 HISTORY OF THE PERIODIC TABLE • Glenn Seaborg • discovered that the transuranium elements that have atomic numbers from 94 to 102, resulting in the redesign of the periodic table • Technically, both the lanthanide and actinide series of elements are to be placed between the alkaline earth metal and the transition metal. • However, by doing this, the periodic table would be too wide. • Thus, the lanthanide and actinide series of elements were placed under the rest of the periodic table. • Dr Seaborg and his colleagues were also responsible for identifying more than 100 isotopes of elements.

46. 4.5 MODERN PERIODIC TABLE The periodic table is a systematic classification of elements whereby elements with the same chemical properties are placed in the same group. • The elements in the periodic table are arranged in rows called the periods and columns which are known as the groups

47. 4.5 MODERN PERIODIC TABLE • Groups • There are 18 groups of elements in the periodic table. • Some of these groups have special names: • (a) Group 1 elements are called alkali metals. • (b) Group 2 elements are called alkaline earth metals. • (c) Group 3 to Group 12 elements are known as transition elements. • (d) Group 17 elements are called halogens. • (e) Group 18 elements are called noble gases. • Each member of a group shows similar chemical properties although their physical properties such as density, melting point and colour show a gradual change when descending the group.

48. 4.5 MODERN PERIODIC TABLE • Periods • There are seven rows from period 1 to period 7. • The elements are arranged horizontally in the ascending order of their proton numbers in the periodic table. • The position of the period of an element in the periodic table is determined by the number of shells occupied with electrons in the atom. • Period 1 has 2 elements only H and He, • Periods 2 and 3 have 8 elements each. • Periods 4 and 5, they have 18 elements each and they are called the long periods. • Period 6 has 32 elements whereas the elements with proton number 58 to 71 are separated and are grouped below the periodic table known as the Lanthanide Series. • Period 7 has 32 elements, the elements with proton number 90 to 103 are grouped below the periodic table known as the Actinide Series.

49. Periodic Classification • Classification as metals and non-metals • (a) Metals – a good conductor of heat and electricity. • (b) Non-metals - a poor conductor of heat and electricity. • (c) Metalloids – a intermediate between metal and non-metal properties

50. ns2np6 ns1 ns2np1 ns2np2 ns2np3 ns2np4 ns2np5 ns2 d10 d1 d5 4f 5f Periodic Classification Ground State Electron Configurations of the Elements The similarity of the outer electron configuration (same type of valence electrons) makes the elements in the same group resemble one another in chemical behavior.