Part 1: The Periodic Table and Physical Properties adapted from Mrs. D. Dogancay

Mendeleev’s Periodic Table Dmitri Mendeleev

Periodic Table Group: elements with same number of valence electrons and therefore similar chemical and physical properties; vertical column of elements (“family”) Period: elements with same outer shell; horizontal row of elements Periodicity: regular variations (or patterns) of properties with increasing atomic weight. Both chemical and physical properties vary in a periodic (repeating pattern) across a period.

Periodic Table

From the IB Data Booklet… 0 or 8 Transition metals Lanthanides Actinides *****page 6 data booklet ********

Another name for “metalloid” is “semi-metal”.

Transition metals alkali metals alkaline earth metals halogens noble gases lanthanides actinides

s, p, d, f blocks

The periodic table is full of repeating patterns. Hence the name periodic table.

PERIODIC TRENDS Properties that have a definite trend as you move through the Periodic Table –Valence Electrons –Effective Nuclear charge –Atomic radii –Ionic radii –Electronegativity –Ionization energy –Electron Affinity –Melting points

The electrons in the outermost electron shell (highest energy level) are called valence electrons. “vale” = Latin to be strong Non Valence electrons are called “core electrons”. Core electrons are relatively stable. Most chemical reactions occur as valence electrons end up in a stable configuration. (full energy level, and to a lesser extent ½ full) Valence electrons

Consider Chlorine: 17 electrons 1s 2 2s 2 2p 6 3s 2 3p 5 Valence Electrons

Consider Chlorine: 17 electrons 1s 2 2s 2 2p 6 3s 2 3p 5 Core ElectronsValence Electrons

Consider Chlorine: 17 electrons 1s 2 2s 2 2p 6 3s 2 3p 5 Core ElectronsValence Electrons Consider Cadmium: 48 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 Valence Electrons

Consider Chlorine: 17 electrons 1s 2 2s 2 2p 6 3s 2 3p 5 Core Electrons7 Valence Electrons Consider Cadmium: 48 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 [Kr]5s 2 4d 10 Valence Electrons 2 Valence Electrons

For the tall groups the column number is the number of valence electrons.

Effective nuclear charge = total protons – core electrons The effective nuclear charge experienced by an atom’s outer electrons ( valence electrons ) increases with the group number across a period s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 1 3p 1 1s 2 2s 2 2p 6 3s 1 3p = = = = + 4

The effective nuclear charge experienced by an atom’s outer electrons ( valence electrons ) remains the same down the group. Effective nuclear charge = total protons – core electrons s 2 2s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s = = = + 1

GENERAL CONCEPTS Many periodic trends can be explained by… 1.How many protons are in the nucleus (effective nuclear charge – positive charge of ion) 2.How far the outer electrons are from the nucleus ( radius- amount of shielding- core electrons ) 3.How stable orbitals p ( d, or f ) are. Ex: p 3 or p 6 4.Energy of orbital ( 3p has more energy than 3s)

+1 Hydrogen Consider Hydrogen v. Helium… electron

+1 +2 Hydrogen Helium Consider Hydrogen v. Helium… Greater (+) Charge in the nucleus produces stronger attraction. electron All orbitals get smaller as protons are added to the nucleus.

Atomic Radii Since the position of the outermost electron can never be known precisely, the atomic radius is usually defined as half the distance between the nuclei of two bonded atoms of the same element. Thus, values not listed in IB data booklet for noble gases. The atomic radius is the distance from the nucleus to the outermost electron. ******page 9 in data booklet******

Atomic Radii Trend: increases down a group WHY??? –The atomic radius gets bigger because electrons are added to energy levels farther away from the nucleus. –Plus, the inner electrons shield the outer electrons from the positive charge (“pull”) of the nucleus; known as the SHIELDING EFFECT

Atomic Radii Trend: decreases across a period WHY??? –As the # of protons in the nucleus increases, the positive charge increases and as a result, the “pull” on the electrons increases.

Ionic Radii ****page 9 of data booklet***** Cations (+) are always smaller than the metal atoms from which they are formed. (fewer electrons than protons & one less shell of e’s) Anions (-) are always larger than the nonmetal atoms from which they are formed. (more electrons than protons) © 2002 Prentice-Hall, Inc.

Ionic Radii Trend: For both cations and anions, radii increases down a group WHY??? –Outer electrons are farther from the nucleus (more shells/ energy levels)

Ionic Radii Trend: For both cations and anions, radii decreases across a period WHY??? –The ions contain the same number of electrons (isoelectronic), but an increasing number of protons, so the ionic radius decreases.

N -3 ( 7 protons) Na + (11 protons) Mg +2 (12 protons) 10 electrons ISOELECTRONIC SERIES Within an isoelectronic series, the ion radii decrease as the atomic number increases ( the electrons are attracted more by the nucleus. N -3 > Mg +2 Na +1 >

Two charges attract each other more if they are _______ and if they are at a _______ distance. Two charges attract each other less if they are _______ and if they are at a _______ distance. Please write : highshort small long

First Ionization Energy Definition: The energy to remove one outer electrons from a gaseous atom. Ionization + + e - g g

Ionization Energy + Work - Energy Doing work against a Coulomb Force

Ionization Energy + Attraction to the nucleus Repulsion from other electrons Work - Energy Doing work against a Coulomb Force Inner electrons tend to “shield” the outer electrons somewhat from the nucleus. This of course is the same effect!

First Ionization Energy K Na Li Ar Ne *****page 8 data booklet ********

First Ionization Energy Trend: decreases down a group. WHY??? –Electrons are in higher energy levels as you move down a group; they are further away from the positive “pull” of the nucleus and therefore easier to remove. –As the distance to the nucleus increases, Coulomb force is reduced. (Remember from Physics the inverse square nature of the force)

The positive value of the ionization energy reminds us that energy must be put into the atom in order to remove the electron. That is to say, the reaction is endothermic. element NaMgAlSiPSClAr # protons Electron arrangement st I.E. (kj mol -1 ) Notice the drop between Mg & Al… evidence of sublevels (s and p)

First Ionization Energy Trend: increases across a period WHY??? –The increasing charge in the nucleus as you move across a period exerts greater “pull” on the electrons; it requires more energy to remove an electron.

First Ionization Energy -WHAT ARE THOSE ELEMENTS? -WHERE DOES THE ELECTRON COME FROM?

Definition: The energy change that occurs when one mole of electrons is added to one mole of gaseous atoms. - + e - gg Electron affinity It can be positive or negative *****page 8 data booklet ********

Electron affinity When the electron affinity is positive, the process is endothermic, the negative ion is less stable than the parent; it will not form spontaneously because there is more repulsion between the electrons. Ex : EA Be = KJ Be + e kJ  Be - Be= [He] 2s 2 The additional electron has to be added to a higher energy sublevel 2p

Electron affinity When the electron affinity is negative, the process is exothermic, the negative ion is more stable than the parent and will form ; the electron is attracted more. Ex: EA of Li = - 60 KJ Li + e  Li kJ Li= [He] 2s 1 The additional electron has to be added to the sublevel 2s where there is space available.

S= Ne 3s 2 3p 4 -> more energy is released, the additional electron is attracted more because S is smaller than P; the nucleus attraction is greater. P = Ne 3s 2 3p 3 -> less energy is released, the addition is less favorable because the electron configuration was a stable p3

Si= Ne 3s 2 3p 2 -> more energy is released, the addition is possible in p 2 and there is an increase in stability in the electron arrangement. P = Ne 3s 2 3p 3 -> less energy is released, the addition is less favorable because the electron configuration was a stable p3

Electronegativity Definition: a relative measure of the attraction that an atom has for a shared pair of electrons when it is covalently bonded to another atom. *****page 8 data booklet ********

Electronegativity Trend: decreases down a group WHY??? –Although the nuclear charge is increasing, the larger size produced by the added energy levels means the electrons are farther away from the nucleus; decreased attraction, so decreased electronegativity; plus, shielding effect

Trend: increases across a period (noble gases excluded!) WHY??? –Nuclear charge is increasing, atomic radius is decreasing; attractive force that the nucleus can exert on another electron increases. Electronegativity

Metals vs Non metals The Metals have lower ionization energies and electronegativities than non-metals: the availability of the valence electrons of the metals explain why they are good conductors of electricity.

Melting Point H He Li Be B C NOF Ne Na Mg Al Si PS Cl Ar K Ca *****page 7 data booklet ********

Melting Point Melting points depend on both… 1.The structure of the element 2.Type of attractive forces holding the atoms together

Melting Point Trend (using period 3 as an example): Elements on the left exhibit metallic bonding (Na, Mg, Al), which increases in strength as the # of valence electrons increases.

Melting Point Silicon in the middle of the period has a macromolecular covalent structure (network) with very strong bonds resulting in a very high melting point.

Melting Point Elements in groups 5, 6 and 7 (P 4, S 8 and Cl 2 ) show simple molecular structures with weak van der Waals’ forces of attraction between molecules (which decrease with molecular size).

Melting Point The noble gases (Ar) exist as single individual atoms with extremely weak forces of attraction between the atoms.

Melting Point Within groups there are also clear trends: In group 1 the m.p. decreases down the group as the atoms become larger and the strength of the metallic bond decreases. elementLiNaKRbCs m.p. (K)

Melting Point Within groups there are also clear trends: In group 7 the van der Waals’ attractive forces between the diatomic molecules increase down the group so the melting points increase. As the molecules get bigger there are obviously more electrons which can move around and set up the temporary dipoles which create these attractions. The stronger intermolecular attractions as the molecules get bigger means that you have to supply more heat energy to turn them into either a liquid or a gas- and so their melting and boiling points rise. At room temperature, chlorine is a gas while iodine is a solid. elementF2F2 Cl 2 Br 2 I2I2 m.p. (K)