Periodic Relationships Among the Elements Chapter 8 Copyright © The McGraw-Hill Companies, Inc.  Permission required for reproduction or display.

When the Elements Were Discovered

Development of the Periodic Table Dimitri Mendeleev (1834 – 1907) Arranged the elements by increasing ATOMIC MASS, he saw a periodic repetition of properties Produced the first PERIODIC TABLE – 1871 The table placed elements with similar properties in the same column Kept “holes” for undiscovered elements, and predicted the properties in advance

Development of the Periodic Table Properties of elements predicted by Mendeleev

Development of the Periodic Table H. G. Moseley (1887-1915) Rearranged the elements by ATOMIC NUMBER This has become the MODERN PERIODIC TABLE

Valence Electrons valence electrons: electrons available to be lost, gained, or shared in the formation of chemical compounds electrons in the outermost energy level electrons that are responsible for reactions Elements in a group have similar properties because they have the same valence electron configuration

All electrons under the highest energy level Inner Core Electrons All electrons under the highest energy level

Valence Electron Configuration Group e- configuration Valence electrons 1 ns1 2 ns2 13 ns2np1 14 ns2np2 15 ns2np3 16 ns2np4 17 ns2np5 18 ns2np6

Charges Of Representative Elements +1 +2 +3 -3 -2 -1 8.2

What atoms are isoelectronic with Ne? an ion that has the same electron configuration Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne] O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne]

Electron Configurations 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

Periodic Table Groups

patterns on the periodic table are called periodic trends PERIODIC LAW When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic (repeating) pattern. patterns on the periodic table are called periodic trends

half the distance from center-center of 2 like atoms Atomic Radius half the distance from center-center of 2 like atoms

Atomic Radii DOWN a Group As you go down a group another energy level is added, the atom size gets larger The number of occupied orbitals between the nucleus and the outermost energy level increases Shielding Effect: reduction of attraction between positive nucleus and outer electrons, outer electrons are not held tight and can move away

Atomic Radius: down group X Na X X X X P X P X P P X P P P P P X P X P X X

Atomic Radius: down group X K X X X X X X X P X X P X P P X X P P P P P X X P X P X X X X

Atomic Radii DOWN a Group DOWN THE GROUP ATOMIC RADIUS INCREASES more energy levels, the larger the size of the atom

Atomic Radii ACROSS a Period Each atom gains one proton and one electron in the same energy level Each added electron is the same distance from the nucleus The positive charge increases and exerts a greater force on the electrons thereby pulling it closer to the nucleus

Effective nuclear charge (Zeff) : “positive charge” felt by an electron. Zeff = Z - s 0 < s < Z (s = shielding constant) Shielding: effect of electrons in the same energy level Zeff  Z – number of inner or core electrons ~Zeff Core Z Radius Na Mg Al Si 11 12 13 14 10 1 2 3 4 186 160 143 132 2s1 would have shielding from 1s2 but 1s2 does not have shielding form 2s1 Outer energy levels have shielding form inner, but its because of the repulsion of inner core electrons with the shielding constant Shielding constant is very small and alters exact charge * For our purposes, close enough to number of valence electrons Within a Period as Zeff increases radius decreases

P e PROTONS REMEMBER! - + are bigger and stronger! electrons are smaller and weaker!

Atomic Radius: across period X X X X X P X P P X P P P P P X P X P X X

Atomic Radii ACROSS a Period ACROSS THE PERIOD ATOMIC RADIUS DECREASES stronger attraction of protons, easier to hold on to the electrons

Atomic Radii

half the distance from center-center of 2 like ions Ionic Radii half the distance from center-center of 2 like ions

Ionic Radius ACROSS the Period Cation: positive ion formed from losing an electron A cation is always smaller than the original atom The more electrons lost the more protons available to attract a smaller number of electrons.

Ionic Radius P X Na + X X P P X P P P P P P P P P

Ionic Radius ACROSS the Period ACROSS THE PERIOD IONIC RADIUS DECREASES greater pull on electrons, the shorter the radius

Ionic Radius ACROSS the Period Anion: negative ion formed from gaining an electron A anion is always larger than the original atom The more electrons gained, the less protons available to attract a larger number of electrons.

Ionic Radius P X F - X X X X P X P P P X P P P P P X P P X X

Ionic Radius ACROSS the Period ACROSS THE PERIOD IONIC RADIUS DECREASES more electrons added, more difficult to keep track of, increasing the size of the atom

Ionic Radius DOWN a Group As you go down a group another energy level is added, increasing the size of the atom. (just like the atomic radius)

Ionic Radius DOWN the Group DOWN THE GROUP IONIC RADIUS INCREASES more energy levels, increase in atom size

8.3

Ionization Energy minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state amount of energy needed to remove an electron from an atom Ion: an atom that has gained or lost electrons if you lose an electron… if you gain an electron… BECOMES POSITIVE! BECOMES NEGATIVE!

Multiple Ionization Energies I1 + X (g) X+(g) + e- I1 first ionization energy I2 + X (g) X2+(g) + e- I2 second ionization energy I3 + X (g) X3+(g) + e- I3 third ionization energy I1 < I2 < I3

Ionization Energy DOWN a Group As you go down a group atoms become larger, electrons are farther from the nucleus and more easily removed The more electrons in an atom between the nucleus and valence shell, the greater the shielding effect

Ionization Energy DOWN a Group DOWN THE GROUP IONIZATION ENERGY DECREASES greater distance from the nucleus, the easier to lose an electron (less energy needed)

Ionization Energy ACROSS a Period As atomic radius decreases there is a greater attraction between protons and electrons. The stronger the attraction, the more energy needed to remove an electron. The more electrons present, the more energy required to remove them all to become STABLE

Ionization Energy ACROSS a Period ACROSS THE PERIOD IONIZATION ENERGY INCREASES more electrons on an energy level, more energy required to remove them all

Filled n=1 shell Filled n=2 shell Filled n=3 shell Filled n=4 shell Filled n=5 shell

: . H F . : : Electronegativity tendency for an atom to attract electrons It is a “tug of war” between the two atoms of a bond . : . H F . : : Which is the more electronegative element?

Electronegativity ACROSS the Period As you go across a period atomic radius decreases because there is a greater attraction between protons and electrons Metals do not attract electrons. Non-metals do attract electrons.

Electronegativity ACROSS the Period ACROSS THE PERIOD ELECTRONEGATIVITY INCREASES stronger the attraction, the easier to add more electrons

Electronegativity DOWN the Group The farther away from the nucleus, the greater the shielding effect The larger the atom, the less likely it is to accept more electrons.

Electronegativity DOWN the Group DOWN THE GROUP ELECTRONEGATIVITY DECREASES farther the distance from the nucleus, more difficult to attract electrons

Increases with ability to attract and hold an electron Electron Affinity the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion. Increases with ability to attract and hold an electron

Electron Affinity A large positive value means the anion is very stable X (g) + e- X-(g) F (g) + e- X-(g) DH = -328 kJ/mol EA = +328 kJ/mol O (g) + e- O-(g) DH = -141 kJ/mol EA = +141 kJ/mol

Electron Affinity ACROSS the Period Across the period the energy released becomes more negative, making electron affinity more positive As the atomic radius decreases and shielding is constant, it is easier to attract an electron. Non-metals want to attract electrons to become stable

Electron Affinity ACROSS the Period ACROSS THE PERIOD ELECTRON AFFINITY INCREASES more electrons and smaller, easily forms anions

Electron Affinity DOWN the Group Down a group, the energy released becomes less negative so electron affinity is smaller The larger the atom, the more difficult to accept electrons Metals want to lose electrons.

Electron Affinity DOWN the Group DOWN THE GROUP ELECTRON AFFINITY DECREASES farther the distance from the nucleus, forms poor anions

Reactivity of Metals Video 1 Reactivity of Metals Video 2 Other Trends Reactivity of Metals Video 1 Reactivity of Metals Video 2

Increasing reactivity METAL REACTIVITY Increasing reactivity

Increasing reactivity NONMETAL REACTIVITY Increasing reactivity