Topic 3 Periodic Trends

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1 Topic 3 Periodic Trends Chapter 06 Trends on the Periodic Table Chapter 07 Relationships between the elements CHEM 10 T03D01

2 How are elements arranged Prior to 1735, only 12 elements were known to man Ag, Cu, Fe, Pb, Sn, Hg, S, C, As, Sb, P, Zn By 1900, 71 new elements (83 total) were known How would they be organized? Mass Physical Properties Chemical Properties Abundance Many tried, the eventual result was the periodic table of elements

3 Attempts at organization 1789 Levoisier s Chemical Textbook Made a table of known elements (O, N, H, etc) He included light and caloric 1817 Dobereiner s Triads Groups of 3 elements had similar properties 1865 Newland s Octaves Classified 56 elements into groups of 8 (Octaves) 1869 Mendeleev s Periodic Table Ordered by relative atomic mass

4 Mendeleev s Periodic Table The most important part about Mendeleev s discovery was that he could predict the properties of unknown elements prior to their discovery!!

5 Periodic Table Describe the arrangement of elements in the periodic table in order of increasing atomic number. The Periodic Law states that when the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties Mendeleev organized all of the elements into one comprehensive table. Elements were arranged in order of increasing mass. Once that was completed elements with similar properties were placed below each other in the same group

6

7 H GROUPS Similarities: The number of electrons in the outer shell. Common reactivity, bonding, chemical and physical properties. Li Be B C N O F Ne Na Mg Al Si P S Cl Ar He K Ca Se Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

8 H METALIC PROPERTIES Similarities: An elements relative ability to conduct energy in the form of heat or electricity. Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Se Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ds the stair Non metals He Metals Metaloids Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

9 H Alkali Metals Non-metals Halogens Metaloids Alkaline Earths Li Be B C N O F Ne Transition Metals Weak/Poor Metals Na Mg Al Si P S Cl Ar He K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Lanthanides Noble Gases Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Actinides

10 Groups vs Periods Distinguish between the terms group and period.(2)

11 Electron Arrangement vs Position Apply the relationship between the electron arrangement of elements and their position in the periodic table up to Z = The electron arrangement repeats, just as the valence electron s do. Once you pass 20 and get into the d-block and f- block electrons, the trend still continues but it makes more sense when represented in the electron configurations of 1s 2 2s 2 2p 6, etc

12 Periodic Trends When looking at periodic trends there are two main properties to compare which help to explain each trend: The Electron Shielding Effect Nuclear Charge You will need to be able to use each of these concepts to describe atomic radius, electronegativity, and ionization energy

13 The Electron Shielding Effect REPULSIONS: If in the same period: Same shell, minimal effect +1 electron, +1 proton B C N slight increase in shielding If in the same group: New shell, more repulsions, LARGE effect +8 electrons, +8 protons B Al Ga large increase in shielding.1

14 Nuclear Charge Effect ATTRACTIONS: If in the same period: Same shell, more E in that shell and P in the nucleus attracted to each other, LARGE effect B C N large increase in nuclear charge If in the same group: New shell, same relative amount of valence E in the shell B Al Ga small increase in nuclear charge +

15 Atomic Radius Down Groups 1 & 7 and across Periods 2&3 Metallic radius Covalent radius Van der Waals radius For metals, the atomic radius is measured by half the distance between two atoms in a metal lattice For non-metals in covalent bonds, the atomic radius is measured by half the distance between the nuclei of two covalently bound atoms For noble gases, the atomic radius is the distance between the nuclei and the outermost electron of an isolated atom

16 The variation of atomic radii in the Alkali Metals Atom Atomic Number Atomic Radius/pm Li Na K Rb Cs Atomic Radius/pm Atomic Radius/pm Fr The variation of atomic radii in the Halogens Atom Atomic Number Atomic Radius/pm F 9 58 Cl Br I At Atomic Radius / pm Atomic Radius / pm

17 Atomic Radius Across Periods 2&3 There is a gradual decrease in atomic radii across periods 2&3. As one more electron and proton are added across: Electrons are added to the same shell and therefore only a slight increase in shielding Additional protons in the nucleus increases the nuclear charge and the electrons in the valence shell are pulled more closely

18 - 3p 10p - 6p Lithium (Li) SAME PERIOD: Increased Repulsions < Increased Attractions Gets SMALLER! p - Carbon (C) Neon (Ne) Argon (Ar) SAME GROUP: Increased Repulsions < Increased Attractions Gets BIGGER! -

19 Other Trends: This concept of size dictated by attractive and repulsive from nuclear charge and electron shielding respectively can be applied to other trends in the periodic table: Ionization Energy Electron Affinity (Electronegativity)

20 Electronegativity and Ionization E Define the terms first ionization energy and electronegativity. The First Ionization Energy is the minimum energy required to remove one mole of electrons from one mole of gaseous atoms (under STP conditions) In general X(g) X + (g) + For example: H(g) H + (g) + ΔH = kj mol -1 It takes 1310 kj of energy to strip away the first and outermost electron from one mole of hydrogen The Second Ionization Energy would be the energy required to take the second electron and is very dependent on the placement in the PT.

21 Trend in 1 st ionization energy Down a group, as atomic radius increases, and increased electron shielding is present, the valence electrons are forgotten Atom Further out from the nucleus Less attraction between protons and valence electrons More electron shielding (repulsions) Requires less energy to remove an electron Atomic Number 1 st I.E. kj mol -1 Li Na K Rb Cs st Ionization Energy (kj/mol) Trend in Ionization Energy for Alkali Metals M.P. (K) Atomic Number

22 1 st Ionization Energy Across Group 3 There is an increase in ionization energy across period 3 Due to increased nuclear charge in comparison to a minor increase in the shielding effect The shells are pulled more closely together and the nucleus has a greater attraction to the valence electrons This increases the ionization energy Some decreases can be explained by viewing the electron configuration Element Na Mg Al Si P S Cl 1 st IE kj mol

23 Electronegativity Electron Affinity The Electronegativity of an atom is the ability or power of an atom in a covalent bond to attract shared pairs of electrons to itself. These values are based on the Pauling scale and range from 4.0 (F) down to 0.7 (Cs,Fr) We will revisit this scale in Bonding Values are pure numbers and simply relative to one another and do not hold units

24 Trend in Electronegativity Values decrease down a group Due to an increased radius and distance between the nucleus and shared pairs of electrons. Nuclear charge is increased but is counteracted by additional electron shielding Values increase across a period As more electrons are added to the valence shell, the nucleus has more of an attraction, the radii decreases, and stability is nearer.

25 Li Na K Li + Na + K + Radii of ionic species Radius increases down a group Radius decreases across a period Cations are smaller as they have an increase effective nuclear charge Anions are bigger as they have a decreased effective nuclear charge F Cl Br F - Cl - Br -

26 Isoelectronic Species Isoelectronic Species are atoms that have the same number of electrons. The Effective Nuclear Charge is the ratio protons to electrons. The greater the ratio, the smaller radius will be relative to it s neutral state Species Na + Mg 2+ Al 3+ Nuc. Charge # electrons Ion Radii (pm) Each are individual sets (of 3) isoelectronic species The two sets cannot be compared as the second set has another complete shell Species P 3- S 2- Cl - Nuc. Charge # electrons Ion Radii (pm)

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