II. Atom. Periodic Table and Trends

1. II. Atom. Periodic Table and Trends Prepared by PhD Halina Falfushynska

2. Rutherford model (planetary model) of atom

3. Draw a diagram showing the location of each part of the atom.

4. Components of Atoms The components that make up the atom are known as sub-atomic particles. Some particles have charges which are integer multiples of the elementary charge.

5. For comparison: • Human hair ~106 atoms wide • Pencil line ~106 atoms wide • HIV virus ~800 atoms wide (~108 atoms total) • E. coli ~1011 atoms total

6. Line Spectra and the Bohr Model • Bohr Model • Since the energy states are quantized, the light emitted from excited atoms must be quantized and appear as line spectra. • After lots of math, Bohr showed that • where n is the principal quantum number (i.e., n = 1, 2, 3, … and nothing else).

7. Line Spectra and the Bohr Model • Bohr Model • We can show that • When ni > nf, energy is emitted. • When nf > ni, energy is absorbed

8. Line Spectra and the Bohr Model Bohr Model

9. Quantum Mechanics and Atomic Orbitals • Schrödinger proposed an equation that contains both wave and particle terms. • Solving the equation leads to wave functions. • The wave function gives the shape of the electronic orbital. [“Shape” really refers to density of electronic charges.] • The square of the wave function, gives the probability of finding the electron ( electron density ).

10. Quantum Mechanics and Atomic Orbitals Solving Schrodinger’s Equation gives rise to ‘Orbitals.’ These orbitals provide the electron density distributed about the nucleus. Orbitals are described by quantum numbers.

11. Quantum Mechanics and Atomic Orbitals • Orbitals and Quantum Numbers • Schrödinger’s equation requires 3 quantum numbers: • Principal Quantum Number, n. This is the same as Bohr’s n. As n becomes larger, the atom becomes larger and the electron is further from the nucleus. ( n = 1 , 2 , 3 , 4 , …. ) • Angular Momentum Quantum Number, . This quantum number depends on the value of n. The values of  begin at 0 and increase to (n - 1). We usually use letters for  (s, p, d and f for  = 0, 1, 2, and 3). Usually we refer to the s, p, d and f-orbitals. • Magnetic Quantum Number, m. This quantum number depends on . The magnetic quantum number has integral values between -  and + . Magnetic quantum numbers give the 3D orientation of each orbital.

12. Quantum Numbers of Wavefuntions

13. Representations of Orbitals The s-Orbitals

14. Representations of Orbitals The p-Orbitals

15. d-orbitals

16. Orbitals and Their Energies Many-Electron Atoms

17. Many-Electron Atoms Electron Spin and the Pauli Exclusion Principle

18. Many-Electron Atoms • Electron Spin and the Pauli Exclusion Principle • Since electron spin is quantized, we define ms = spin quantum number =  ½. • Pauli’s Exclusions Principle:no two electrons can have the same set of 4 quantum numbers. • Therefore, two electrons in the same orbital must have opposite spins.

19. Orbitals CD Figure 6.27

21. Atomic and Mass Numbers, Isotopes • Using the periodic table, the atomic number (number of protons; used to identify the element), elemental symbol, and atomic weight can be identified.

22. Nuclides • Atoms may exist as more than one isotope – atoms with the same number of protons but have a different number of neutrons. The mass number is used to distinguish isotopes. • One of two or more atoms whose nuclei have the same number of neutrons but different numbers of protons are called isotones. • One of two or more atoms or elements having the same atomic weights or mass numbers but different atomic numbers are called isobars

23. Average atomic weight • The average atomic weight is a weighted average of naturally occurring isotopes and fractional abundance in nature. Example: Carbon has two appreciably present isotopes, 12C and 13C. Respectively, the abundances are 98.9% and 1.1%. • average atomic weight = (0.989)(12 amu) + (0.011)(13 amu) = 12.011 amu (on periodic table) av. atomic wt. = (mass isotope A)(% A) + (mass isotope B)(% B)

24. Molecules and Molecular Compounds • Sharing electrons molecular (covalent) compound • Trading electrons  ionic compound

25. Charges • A cation is formed when more protons than electrons are present in an atom or molecule • An anion is formed when more electrons than protons are present in an atom or molecule.

26. Examples O Na 2- F + 18 23 19 8 11 9 8 protons 10 neutrons 10 electrons 9 protons 10 neutrons 9 electrons 11 protons 12 neutrons 10 electrons

27. Predicting Ionic Charges • Metals form cations. • Nonmetals form anions.

28. Naming Inorganic Compounds • Positive ions: • • Cations formed from metal (main group or transition) atoms have the same name. • • If a metal can form different cations, the positive charge is indicated by a Roman numeral in parenthesis following the name of the metal. Older names using –ous and –ic are still seen but their use is fading. • • Cations formed from nonmetals have names that end in –ium (i.e. hydronium ion and ammonium ion).

29. The Periodic Law says: • When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties. • Horizontal rows = periods • There are 7 periods • Vertical column = group (or family) • Similar physical & chemical prop. • Identified by number & letter (IA, IIA)

31. Areas of the periodic table • Three classes of elements are: • Metals: electrical conductors, have luster, ductile, malleable • Nonmetals: generally brittle and non-lustrous, poor conductors of heat and electricity • Metalloids:border the line-2 sides • Properties are intermediate between metals and nonmetals

32. Electron Configurations in Groups • Elements can be sorted into 4 different groupings based on their electron configurations: • Noble gases • Representative elements • Transition metals • Inner transition metals

33. Classify elements based on electron configuration. • Group IA – alkali metals • Forms a “base” (or alkali) when reacting with water (not just dissolved!) • Group 2A – alkaline earth metals • Also form bases with water; do not dissolve well, hence “earth metals” • Group 7A – halogens • Means “salt-forming”

34. Electron Configurations in Groups • Noble gasesare the elements in Group 8A(also called Group18 or 0) • Previously called “inert gases” because they rarely take part in a reaction; very stable= don’t react • Noble gases have an electron configuration that has the outer s and p sublevels completely full

35. Electron Configurations in Groups • Representative Elementsare in Groups 1A through 7A • Display wide range of properties, thus a good “representative” • Some are metals, or nonmetals, or metalloids; some are solid, others are gases or liquids • Their outer s and p electron configurations are NOT filled

36. Electron Configurations in Groups • Transition metalsare in the “B” columns of the periodic table • Electron configuration has the outer s sublevel full, and is now filling the “d” sublevel • A “transition” between the metal area and the nonmetal area • Examples are gold, copper, silver

37. Electron Configurations in Groups • Inner Transition Metals are located below the main body of the table, in two horizontal rows • Electron configuration has the outer s sublevel full, and is now filling the “f” sublevel • Formerly called “rare-earth” elements, but this is not true because some are very abundant

39. H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 Do you notice any similarity in these configurations of the alkali metals? 1s1 1s22s1 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p63s23p64s23d104p65s1 1s22s22p63s23p64s23d104p65s24d10 5p66s1 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1

40. ALL Periodic Table Trends • Influenced by three factors: 1. Energy Level • Higher energy levels are further away from the nucleus. 2. Charge on nucleus (# protons) • More charge pulls electrons in closer. (+ and – attract each other) • 3. Shielding effect (blocking effect?)

41. What do they influence? • Energy levelsand Shielding have an effect on the GROUP (  ) • Nuclear chargehas an effect on a PERIOD (  )

42. #1. Atomic Size - Group trends H • As we increase the atomic number (or go down a group). • each atom has another energy level, • so the atoms get bigger. Li Na K Rb

43. #1. Atomic Size - Period Trends • Going from left to right across a period, the size gets smaller. • Electrons are in the same energy level. • But, there is more nuclear charge. • Outermost electrons are pulled closer. Na Mg Al Si P S Cl Ar

45. Ions • Some compounds are composed of particles called “ions” • An ion is an atom (or group of atoms) that has a positive or negative charge • Atoms are neutral because the number of protons equals electrons • Positive and negative ions are formed when electrons are transferred(lost or gained) between atoms

46. Ions • Metals tend to LOSE electrons, from their outer energy level, and thus a positively charged particle is formed = “cation”. Cations are smaller than the atom they came from • Nonmetals tend to GAINone or more electrons. Negative ions are called “anions”. Anions are bigger than the atom they came from

47. Ion Group trends • Each step down a group is adding an energy level • Ions therefore get biggeras you go down, because of the additional energy level. Li1+ Na1+ K1+ Rb1+ Cs1+

48. Ion Period Trends • Across the period from left to right, the nuclear charge increases - so they get smaller. • Notice the energy level changesbetween anions and cations. (more protons would pull the same # of electrons in closer) O2- F1- N3- B3+ Li1+ Be2+ C4+

50. #2. Trends in Ionization Energy • Ionization energy is the amount of energy required to completely remove an electron(from a gaseous atom). • Removing one electron makes a 1+ ion. • The energy required to remove only the first electron is called the first ionization energy.