# Coincidence? I Think Not! 0 As you have realized, the Periodic Table provides a great deal more information than just atomic number and atomic mass!

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2 Coincidence? I Think Not! 0 As you have realized, the Periodic Table provides a great deal more information than just atomic number and atomic mass! 0 Each period (row) corresponds to an energy level 0 Each atomic orbital (s, p, d, f) is represented as a block 0 This is because elements on the Periodic Table are arranged by increasing atomic number 0 As a result, a regular and repeating pattern of chemical and physical properties emerges 0 Called periodic law Blue = s block Red n = 1 Orange n = 2 Yellow n = 3 Green n = 4 Blue n = 5 Indigo n = 6 Violet n = 7

3 More on the Periodic Law 0 Atoms with similar properties appear in groups or families (vertical columns) on the periodic table 0 They are similar because they all have the same number of valence (outer shell) electrons, which governs their chemical behavior 0 Remember, valence electrons are electrons in the highestnumbered s- and p- orbitals! 0 Elements of the same period (horizontal row) have the same number of energy levels 0 As you move across a period, the number of electrons and protons increases, leading to increase in atomic number 0 Elements within the same period do not generally show similarity in properties, except d-block and f-block (lanthanides) elements

4 Trends in the Periodic Table 0 Due to this arrangement, patterns of chemical and physical properties emerge 0 Called trends 0 These trends occur due to the electronic configurations of elements!

5 So, What are These Trends?

6 Trend #1 - Atomic Radius 0 Defined as the distance between the nucleus and the outer edge of the electron cloud 0 Since an electron cloud s edge is difficult to define, scientists use the covalent radius, or half the distance between the nuclei of 2 bonded atoms 0 Atomic radii are usually measured in picometers (pm) or angstroms (Å). 0 An angstrom is 1 x m

7 Example Atomic Radius 0 Two bromine atoms bonded together are 2.86 angstroms apart. So, the radius of each atom is 1.43 Å Å 1.43 Å 1.43 Å

8 Atomic Radius Trend -Across a Period 0 Atomic radius DECREASES from left to right across a period 0 Why? 0 As you move across a period, a proton AND electron are added, as well as 1 or 2 neutrons 0 The result is a more positive nucleus and a more negative electron cloud 0 The outer-shell electrons experience an increased attraction to the nucleus and the electron cloud gets pulled in, making atoms smaller 0 The amount of positive charge perceived by an electron from the nucleus is called the nuclear effective charge (Z eff ) 0 Z eff increases across a period

9 Atomic Radius Trend Down a Group 0 Atomic radius INCREASES down a group 0 Why? 0 Number of energy levels (n) increases so, orbitals are larger 0 The bigger the distance over which the nucleus must pull reduces the attraction for electrons 0 Furthermore, outer-shell electrons are shielded from the full nuclear positive charge by the inner-shell electrons Atomic Radius Animation

11 Trend #2 - Ionic Radius 0 Defined as the distance from the nucleus to the other edge of the electron cloud in a charged ion 0 Metal atoms can lose valence electrons to form positive ions called cations 0 Nonmetal atoms can gain valence electrons to form negative ions called anions

12 Ionic Radius Trend Across a Period and Down a Group 0 Same radii trends apply once you divide the table into metal (cation) and non-metal (anion) sections 0 Cation radii DECREASES from left to right with only minor changes in the transition metals 0 Anion radii are larger than cations and DECREASES from left to right 0 As electrons are added, the p+/e- ratio decreases and the electrons are not as closely held 0 Increased electrons and electron repulsions also play a role in expanding the electron cloud 0 Ionic radii increases down all groups because of the additional energy levels (n)

13 Sodium Cation Formation 1 valence electron Effective nuclear charge on remaining electrons increases! 11p + Remaining e- are pulled in closer to the nucleus - ionic size decreases! Valence e- lost in ion formation Result - a smaller sodium cation, Na +

14 Chlorine Anion Formation 7 valence e- A chloride ion is produced - it is larger than the original atom! 17p + One e- is added to the outer shell Effective nuclear charge is reduced and the e- cloud expands!

15 K +1 Cs + is larger than K + Br - is bigger than Cl - When going down a group, the ions get bigger Cs +1 Cs +1 This is true for cations & anions!

16 GOING ACROSS A PERIOD CATIONS get smaller ANIONS get smaller too

18 Trend #3 - Ionization Energy (IE) 0 Defined as the energy needed to remove an electron from a gaseous atom or ion 0 In other words if an electron is given enough energy (in the form of a photon) to overcome the effective nuclear charge in a gaseous atom or ion, it can leave the atom completely and become ionized or charged 0 This is because the number of protons and electrons is no longer equal 0 IE is measured in kj 0 The larger the I.E, the harder it is to remove the electron 0 This process is ALWAYS endothermic 0 This process is also a stepwise process 0 There are first, second, third, etc. ionization energies 0 First ionization energy is the energy required to remove the highestenergy electron (one bound least tightly) 0 Second ionization energy is the energy required to remove an additional electron from the mole of ions (+1 cations into +2 cations) 0 A large jump in ionization energy occurs when a CORE e - is removed

19 Example Ionization Energies of Magnesium 1st I.E. 2nd I.E. 3rd I.E. 736 kj 1,445 kj 7,730 kj Core e -

20 First Ionization Energy Trend Across a Period 0 FIRST ionization energy INCREASES left to right across a period 0 First ionization energy and atomic radius are inversely proportional to each other 0 Why? 0 Effective nuclear charge, Z eff, increases the attraction of the nucleus and therefore, pulls the electron cloud closer to the nucleus 0 Holds electrons more tightly harder to remove

21 Exceptions to the I.E Trend 0 Ionization energies do not increase smoothly across the periods in the periodic table 0 In general: - It is easier to remove an electron if it results in the formation of a filled or half-filled subshell - It is harder to It is harder to remove an electron from a filled or halffilled subshell

22 First Ionization Energy Trend Down a Group 0 FIRST ionization energy DECREASES down a group 0 Why? 0 The larger the atom is, the easier its electrons are to remove 0 Increased number of energy levels (n) increases the distance over which the nucleus must pull and therefore, reduces the attraction for electrons 0 Also, full energy levels provide some shielding between the nucleus and valence electrons

23 Need More Clarification? Ionization Energy Video

24 3-D Graph of First IE

25 Summary of First IE Trend

26 Trend #4 - Electronegativity 0 Defined as a measurement of the attraction of an atom of the pair of outer shell electrons in a covalent bond with another atom 0 In other words, electronegativity is a measure of an atom s attraction for another atom s electrons 0 It is an arbitrary scale that ranges from 0 to = no electronegativity 0 4 = high electronegativity 0 All electronegativity values are relative to fluorine, the most electronegative element

27 Electronegativity Trend Across a Period 0 Generally, metals are electron givers and have low electronegativities 0 Nonmetals are electron takers and have high electronegativities 0 What do you think about the electronegativity of noble gases? 0 So, electronegativity INCREASES left to right across a period 0 Why? 0 Effective nuclear charge, Z eff, increases the attraction of the nucleus and therefore, strengthens the attraction for electrons 0 Holds electrons more tightly

28 Electronegativity Trend Down a Group 0 Electronegavity DECREASES down a group 0 Why? 0 Increased number of energy levels (n) increases the distance over which the nucleus must pull and therefore, reduces the attraction for electrons 0 Also, full energy levels provide some shielding between the nucleus and valence electrons

29 Summary of Electronegativity Trend 0

30 A General Justification of the Trends 0 Across a period, electrons of the outermost shell experience increased nuclear attraction due to increase in atomic number 0 Called nuclear effective charge (Z eff ) the amount of positive charge perceived by an electron from the nucleus 0 Down a group, nuclear attraction decreases due to increased distance from the nucleus and the shielding effect 0 Electrons intervening between the nucleus and an outer electron are said to shield or screen the outer electron from the nucleus 0 As a result, the outer electron does not experience the full nuclear charge 0 Only full energy levels, not full orbitals (s, p, d, f), are of concern in a shielding argument

31 Summary of All Periodic Trends

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