Chapter 6 The Periodic Table Jennie L Borders

Chapter 6 – The Periodic Table Jennie L. Borders

Section 6. 1 – Organizing the Periodic Table • Chemists used the properties of elements to sort them into groups.

Mendeleev • Mendeleev is credited with creating the first useful periodic table. • He arranged the elements in order of increasing atomic mass. • He also put elements with similar properties in the same group.

Mendeleev • When he finished, there were blanks in his periodic table. • Since he arranged his periodic table based on properties, he predicted the properties of elements that had not been discovered. • When the elements were discovered, his predictions were right.

Modern Periodic Table • The modern periodic table is arranged in order of increasing atomic number. • Elements in the same group have similar properties. • Elements in the same period have a repeating set of properties. This is referred to as the periodic law.

Metals, Nonmetals, and Metalloids • The periodic table can be broken up into metals, nonmetals, and metalloids.

Metals • Properties of metals include: – Good conductors – Shiny – Solid (except mercury) – Ductile – can be pulled into wires – Malleable – can be hammered into sheets – Low ionization energies – energy needed to remove an electron – Form positive ions – Tend to be oxidized (lose electrons)

Metal Oxides • Most metal oxides are basic. • Bases produce OH- ions when dissolved in water. Na 2 O + H 2 O 2 Na. OH Ca. O + H 2 O Ca(OH)2

Nonmetals • Properties of nonmetals include: – Tend to be gases – Poor conductors (except carbon) – Brittle – Dull – High electron affinities – the energy change that occurs when an atom gains an electron – Form negative ions – Tend to be reduced (gain electrons)

Nonmetal Oxides • Most nonmetal oxides are acidic. • Acids produce H+ ions when dissolved in water. CO 2 + H 2 O H 2 CO 3 P 4 O 10 + 6 H 2 O 4 H 3 PO 4

Metalloids • Metalloids generally have some of the properties of metals and nonmetals.

Sample Exercise • Which two of the following elements would you expect to show the greatest similarity in chemical and physical properties: B, Ca, F, He, Mg, P?

Practice Exercise • Locate sodium and bromine on the periodic table. Give the atomic number of each, and label each a metal, metalloid, or nonmetal.

Section 6. 1 Assessment 1. What property did Mendeleev use to organize his periodic table? 2. How are elements arranged in the modern periodic table? 3. Name three broad classes of elements. 4. Which of these sets of elements have similar physical and chemical properties ? a. oxygen, nitrogen, carbon, boron b. strontium, magnesium, calcium, beryllium c. nitrogen, neon, nickel, niobium

Section 6. 1 Assessment 5. Identify each element as a metal, metalloid, or nonmetal. a. gold b. silicon c. sulfur d. barium 6. Name two elements that have properties similar to those of the element sodium.

Section 6. 2 – Classifying the Elements La Ac

Group Trends for the Active Metals • The alkali metals (group 1) and the alkaline earth metals (group 2) are considered the active metals.

Properties of Alkali Metals • • Soft metallic solids with low melting points When bonded with hydrogen, the hydrogen has a -1 charge (hydride). Ex: Li. H React vigorously with water When bonded with oxygen, they can form the following: a. Oxide (O 2 -) = Li 2 O b. Peroxide (O 22 -) = Na 2 O 2 c. Superoxide (O 2 -) = KO 2

Properties of Alkaline Earth Metals • Harder and more dense than alkali metals • Higher melting point than alkali metals • Less reactive with water than alkali metals alkali metal alkaline earth metal

Group Trends for Selected Nonmetals • Hydrogen is in the alkali metal group even though it is a nonmetal. It can be metallic under extremely high pressures. • It belongs in group 1 because it has 1 valence electron and can form a +1 charge. • It belongs in group 17 because it is a nonmetal, it can form a -1 charge, and it only needs 1 more electron to achieve noble gas configuration.

Nonmetal Groups • When going down a group of nonmetals, the elements go from nonmetallic to metallic in nature.

Group 17 - Halogens • • • Fluorine is a pale yellow gas. Chlorine is a greenish-yellow gas. Bromine is a reddish-brown liquid. Iodine is a gray/black solid that forms purple vapors. Halogens are very reactive.

Group 18 – Noble Gases • Noble gases have full outer energy levels, so they are very unreactive. • They rarely form compounds, but compounds have been formed with xenon, krypton, and argon. Most of these compounds contain fluorine, since it is highly reactive.

Electron Configuration in Groups • Elements in the same group have similar properties because they have similar electron configurations.

Sample Exercise • What is the characteristic valence electron configuration of the group 17 elements, the halogens?

Practice Exercise • Which group of elements is characterized by an ns 2 np 2 electron configuration?

Section 6. 2 Assessment 1. Into what four classes can elements be sorted based on their electron configuration? 2. Why do the elements potassium and sodium have similar chemical properties? 3. Which of the following elements are transition metals: Cu, Sr, Cd, Au, Al, Ge, Co? 4. How many electrons are in the highest occupied energy level of a Group 15 element?

Section 6. 3 – Periodic Trends • Periodic trends are trends that occur as you move down or across the periodic table.

Effective Nuclear Charge (Zeff) • Since electrons are negatively charged, they are attracted to nuclei, which are positively charged. • Many of the properties of atoms depend on their electron configurations and on how strongly their outer electrons are attracted to the nucleus.

Electrons and the Nucleus • The force of attraction between an electron and the nucleus is based on Coulomb’s law and depends on two amounts: 1. The magnitude of the net nuclear charge acting on the electron 2. The average distance between the nucleus and the electron

Effective Nuclear Charge • Effective nuclear charge is the actual nuclear charge minus the shielding of the core electrons, so Zeff is always less than actual nuclear charge. Zeff = Z – S Zeff = effective nuclear charge Z = actual nuclear charge (atomic number) S = screening constant (number of core electrons)

Sample Exercise • Calculate the Zeff for Ar and for Kr. • Which will experience the greater effective nuclear charge in the n=3 subshell? • Which n=3 electrons will be closer to the nucleus?

Practice Exercise • Calculate the effective nuclear charge for Na and for Br.

Zeff Trend Across • The effective nuclear charge increases as we move across a row (period) on the table. • Although the number of core electrons stays the same as we move across, the actual nuclear charge increases. • Ex: Na = 1+ (Zeff = 11 – 10) and Mg = 2+ (Zeff = 12 – 10)

Zeff Trend Down • Going down a column, the effective nuclear charge increases slightly. • Actual nuclear charge is increasing as you move down a group, but the larger electron cores are less able the screen the valence electrons. • However, our simple Zeff calculations make it seem like it stays the same. • Ex: F = 7+ (Zeff = 9 – 2) • Ex: Cl = 7+ (Zeff = 17 – 10)

Atomic Radius/Size • Atomic Radius – the radius of an atom. • In general, the atomic radius increases as you move down a group and decreases as you move across a period. • The trends going across do not change as much when you move across the d and f sublevels

Atomic Radius • The atomic radius increases going down a group because larger energy levels are added with each row. • The atomic radius decreases going across a period because electrons are added to the same energy level, but protons are added to the nucleus (Zeff increases) which pull the electron in closer.

Sample Exercise • Arrange the following atoms in order of increasing size: P, S, As, and Se.

Practice Exercise • Arrange the following atoms in order of increasing atomic radius: Na, Be, Mg.

Ions • An ion is an atom with a charge. An atom has a charge when it gains or loses electrons. • An anion is a negative ion (gains electrons). • A cation is a positive ion (loses electrons).

Charges • You can tell the charge of an element based on which group it is in on the periodic table (except for transition metals).

Ionization Energy • Ionization energy is the energy needed to remove an electron from an atom. • In general, ionization energy decreases as you move down a group and increases as you move across a period.

Ionization Energy • Ionization energy decreases as you move down a group because larger energy levels are added which are farther from the nucleus. Since the electrons are far from the nucleus, it takes less energy to remove one. • Ionization energy increases as you move across a period because the nucleus gets stronger (Zeff increases), so it takes more energy to remove an electron.

Sample Exercise • Arrange the following atoms in order of increasing first ionization energy: Ne, Na, P, Ar, K.

Practice Exercise • Which has the lowest first ionization energy B, Al, C, or Si? • Which has the highest first ionization energy?

Ionization Energies • Notice the values for a given element increase as successive electrons are removed: I 1<I 2<I 3, and so forth. • This trend exists because each extra electron is being removed from an increasingly positive ion.

Ionization Energies • There is a sharp increase in the ionization energy that occurs when an inner-shell electron is removed. • This is the reason that only valence electrons are involved in bonding.

Sample Exercise • Based on their locations on the periodic table, does sodium, calcium, or sulfur have the highest second ionization energy?

Practice Exercise • Which will have the greater third ionization energy, Ca or S?

Ionic Size • Ionic radius is the radius of an ion. • Cations are smaller than the parent atom. • Anions are larger than the parent atom.

Ionic Size • In general, ionic size increases as you move down a group because larger energy levels are added.

Ionic Size • Ionic size generally decreases across the cations, then increases as you move to the anions. As you move across the anions the size decreases again. This is due to the increased strength of the nucleus and the loss or gain of electrons.

Sample Exercise • Arrange these atoms and ions in order of decreasing size: Mg 2+, Ca.

Practice Exercise • Which of the following atoms and ions is largest: S 2 -, S, O 2 -?

Isoelectronic Series • An isoelectronic series is a group of ions all containing the same number of electrons. • In an isoelectronic series we can list the members in order of increasing atomic number and the radius will decrease. Increasing atomic number O 2 - F- Na+ Mg 2+ Al 3+ Decreasing atomic radius

Isoelectronic Series

Sample Exercise • Arrange the ions K+, Cl-, Ca 2+, and S 2 - in order of decreasing size.

Practice Exercise • Which of the following ions is largest, Rb+, Sr 2+, or Y 3+?

Electronegativity • Electronegativity is the ability of an atom to attract more electrons. • In general, electronegativity decreases as you move down a group and increases as you move across a period.

Electronegativity • Electronegativity decreases as you move down a group because larger energy levels are added that are farther from the nucleus so the atom cannot attract electrons as well. • Electronegativity increases as you move across a period because the nucleus is stronger (Zeff increases) and can attract more electrons.

Electron Affinities • The energy change that occurs when an electron is added to a gaseous atom is called the electron affinity. It measures the attraction of the atom for the added electron. • For most atoms, energy is released when an electron is added.

Electron Affinity Trend Across • Electron affinity does not have a consistent trend going across a period. It does overall increase as you move across. • This trend occurs due to the increased Zeff and smaller atomic radius.

Electron Affinity Trend Down • The electron affinity slightly decreases as you move down a group. • This trend is due to the larger energy levels that are farther from the nucleus.

Summary of Trends

Section 6. 3 Assessment 1. How does atomic size change within groups and across periods? 2. When do ions form? 3. What happens to first ionization energy within groups and across periods? 4. Compare the size of ions to the size of the atoms from which they form. 5. How does electronegativity vary within groups and across periods?

Section 6. 3 Assessment 6. Arrange these elements in order of decreasing atomic size: sulfur, chlorine, aluminum, and sodium. Does your arrangement demonstrate a periodic trend or a group trend? 7. Which element is each pair has the larger first ionization energy? a. sodium, potassium b. magnesium, phosphorus

Sample Integrative Exercise • The element bismuth (Bi, atomic number 83) is the heaviest member of group 15. A salt of the element, bismuth subsalicylate, is the active ingredient in Pepto. Bismol, an over-the-counter medication for gastric distress. a. The covalent atomic radii of thallium (Tl) and lead o o (Pb) are 1. 48 A and 1. 47 A, respectively. Using these values, predict the atomic radius of the element bismuth. Explain your answer.

Sample Integrative Exercise b. What accounts for the general increase in atomic radius going down the group 15 elements? c. Another major use of bismuth has been as an ingredient in low-melting metal alloys. The element itself is a brittle crystalline solid. How do these characteristics fit with the fact that bismuth is in the same periodic group with such nonmetallic elements as nitrogen and phosphorus?

Sample Integrative Exercise d. 209 Bi is this heaviest stable isotope of any element. How many protons and neutrons are present in the nucleus?