PERIODICITY

1. PERIODICITY Chelsea Greenberg And Spenser Jacobson

2. KEY TERMS • Valence electrons- the electrons contained in the outermost, or valence, electron shell of an atom; these electrons determine how the atom can bond. • Ionizationenergy- the energy required to remove a valence electron from an atom • The first ionization energy is the energy required to remove the first valence electron • Effective nuclear charge- electrons are attracted to the protons in an atom and electrons are repelled by other electrons in an atom, these partials charges are the basis of periodic trends • Paramagnetic-some of the element’s electrons are unpaired and thus effected by magnetic fields • Diagmagnetic- all electron are paired, and their subshells are complete • Stability- Atoms prefer to have complete shells if possible in order to be stable • Heat of reaction- the amount of heat absorbed or released in a reaction • When bonds are formed, energy is released • When bonds are broken, energy is absorbed • Physical property-can be observed and measured without the composition of matter changing. They are used to observe and describe matter. • Chemical property-describes an element’s potential to undergo some chemical change of reaction by virtue of its composition.

3. The covalent radii decrease regularly from Li to F, whereas the first ionization energies do not. For the ionization energies, show how currently accepted theoretical concepts can be used to explain the general trend and the two discontinuities. The atomic radius decreases from left to right because there are more protons per atom which results on a stronger pull on the electrons. Because of the stronger attraction, and greater effective nuclear charge, it is increasingly difficult to remove an electron and the ionization energy of an atom increases from left to right across the periodic table. There is more energy required to remove an electron from an atom on the right of a period than from one on the left. However, there are exceptions to this trend.

4. First Ionization Energy (kc/mole) Li 124 Be 215 B 191 C 260 N 336 O 314 F 402 There are exceptions to the general trend. Boron has a lower first ionization energy than Beryllium, meaning it is easier to remove the outermost valence electron from Boron than from Beryllium. This can be accounted for based on the element’s ground state electron configuration The covalent radii decrease regularly from Li to F, whereas the first ionization energies do not. For the ionization energies, show how currently accepted theoretical concepts can be used to explain the general trend and the two discontinuities. The ground state configuration of Boron has an electron in a higher energy orbital than Beryllium. This electron is in 2p and is more easily ionized than the electron from 2s. Therefore, although Boron is right of Beryllium, Boron has a lower first ionization energy.

5. The covalent radii decrease regularly from Li to F, whereas the first ionization energies do not. For the ionization energies, show how currently accepted theoretical concepts can be used to explain the general trend and the two discontinuities. First Ionization Energy (kc/mole) Li 124 Be 215 B 191 C 260 N 336 O 314 F 402 There is another exception to the ionization trend. The first ionization energy for Oxygen is less than that of Nitrogren. This is because the electron to be ionized in Oxygen is a paired electron in the 2p orbital. Nitrogen the atom is more stable and less easily ionized because it is parametric. Because Oxygen has only one paired set of electrons, it is less symmetric and one electron is easily removed to result in greater stability.

6. The reaction of a metal with chlorine proceeds as indicated below. Indicate, with reasons for your answers, the effect of the following factors on the heat of reaction for this reaction. M(s) + Cl2(g) MCl2(g) • A large radius versus a small radius for M2+ The resulting molecule, MCl2 will be bigger when M is an atom with a large radius. For exampled Barium has a larger radius that Beryllium and Barium Chloride will be a bigger molecule than Beryllium Chloride.The smaller the radius, the closer the atoms can bond and the smaller the resulting molecule. Because the atoms cannot get as close when the atom has a large radius, it is more difficult to form a bond. This means that more energy is released, and the reaction is more exothermic. The larger the radius of M, the more heat that is released in the reaction.

7. The reaction of a metal with chlorine proceeds as indicated below. Indicate, with reasons for your answers, the effect of the following factors on the heat of reaction for this reaction. M(s) + Cl2(g) MCl2(g) B) A high ionization energy versus a low ionization energy for M. When an atom with a higher ionization energy is used, the reaction is more exothermic. Beryllium has a higher ionization energy than Barium. This means that is harder for Beryllium to give up an electron and is less likely to bond. When Beryllium does bond it requires more energy than when Barium bonds and results in a greater heat of formation. When bonds are formed, energy is released. Atoms with a higher ionization energy makes it more difficult for bonds to form and a greater heat of reaction for the reaction above.

11. Discuss some differences in physical and chemical properties of metals and nonmetals. What characteristic of the electronic configuration of atoms distinguishes metals from nonmetals. On the basis of this characteristic explain why there are many more metals than nonmetals. METALS-PHYSICAL PROPERTIES • Malleable-can be hammered or beaten without breaking because their metallic bonds are very strong. Metals can be made into foils • Ductile- can be melted and drawn into thin wires by being stretched. The atoms can stretch and not separate because of the strong metallic bonds • Conductors- metals conduct electricity and heat because of the excess electrons that move around the metallic bonds. • Luster- metals can be highly polished and they shine. Because of the electronic structure of metals, they are able to reflect incident light. • Tensile- metals are very strong because they are ductile and malleable. Their bonds do not break easily as the electrons are shared over an array of metal atoms. • Physical state- metals are solid at room temperature (except for mercury which is a liquid). Metals also have a high density. All metals are hard materials except sodium and potassium which are soft metals. • Brittleness- metals are not brittle except zinc. They do not break easily before of their metallic bonds. • Meting and Boiling Points- all metals except sodium and potassium have high melting and boiling points. • Solubility- Pure metals are insoluble in solvents like water or in any organic solvent. Metals can be dissolved only in acids. • Sonorous- Metals make a characteristic sound when hit with an object. The sonorousness of metals depends on the temperature and density.

12. Discuss some differences in physical and chemical properties of metals and nonmetals. What characteristic of the electronic configuration of atoms distinguishes metals from nonmetals. On the basis of this characteristic explain why there are many more metals than nonmetals. Metals- Chemical Properties • Electropositive- Atoms of metal elements have 1 to 3 electrons in the outermost shells. The atoms can therefore loose electrons to form compound and when they do they become electropositive. • Reaction with Oxygen- metals react with oxygen to form metal oxides which are basic. Metal oxide bonds are ionic and dissolve in water to form alkaline solutions. • Reaction with Water- All metals react with water to form metal oxides and metal hydroxides and hydrogen gas is released. • Dilute Acids- Most metals react with dilute acids by replacing hydrogen and forming a salt and hydrogen gas is released. The rate of reactions depends on the reactivity of the metal. Na reacts vigorously with dilute acids, whereas Zn or Fe react very slowly with dilute acids. • Salt Solutions-reactivity of metals determine which salt is formed when a metal is added to a salt solution. More reactive metal displaces the less reactive metal in the metal salt. These are called displacement reactions. • Chlorides- Metals have excess electrons and Chlorine needs one electron to complete its orbit. Thus metals react very easily with Chlorine. All chlorides are electrovalent or ionic compounds. • Hydrogen- Na, K, Mg, and Ca react easily with H to form metal hydrides. Other metals have to be heated to react with hydrogen. • Reducing Agents- Metals have excess unbound electrons and are able to give these electrons easily. In chemistry this property of a substance to donate electrons is known as reducing property and the substance is called a reducing agent. For example Na gives off an electron to any other substance which is forming a bond with itself.