The Periodic Table

The Modern Periodic Table u The modern periodic table is based on the atomic numbers of the elements.

The Modern Periodic Table u The periodic table is arranged in order of increasing atomic number. u The physical and chemical properties of the elements repeat in a regular pattern when they are arranged in order of increasing atomic number.

u Elements in vertical columns showed similar properties. The Periodic Table For example, alkaline earth metals have high melting points and low density and are silver in color, ductile, and malleable.

Periodicity u This repeated pattern is an example of periodicity in the properties of elements. u Periodicity is the tendency to recur at regular intervals.

The Periodic Table u On the periodic table a period, sometimes also called a series, consists of the elements in a horizontal row.

The Periodic Table u There are 7 periods in the table.

The Periodic Table u A group, sometimes also called a family, consists of the elements in a vertical column.

Metals, Nonmetals, and Metalloids u There are 3 main regions of the table: metals, nonmetals and metalloids.

u The metals are in blue. Metals

Metals u Metals are elements that have luster (are shiny), conduct heat and electricity, and usually bend without breaking (malleable). u Metals are also ductile (can be drawn out into a wire).

Metals u Most metals have one, two, or three valence electrons. u Metals tend to lose electrons in order to achieve the stability of a filled octet.

Metals u All metals except mercury are solids at room temperature; in fact, most have extremely high melting points.

Metal Reactivity u A metal’s reactivity is its ability to react with another substance. u Reactivity for metals increases as you go down a group and left across a period.

Metal Reactivity

1. Consult the “Activity Series of Metals” in the Chemistry Reference Tables to determine the more active metal. (manganese) a) cobalt (Co) or manganese (Mn) b) barium (Ba) or sodium (Na) (barium)

u The nonmetals are in yellow. Nonmetals

Nonmetals u Although the majority of the elements in the periodic table are metals, many nonmetals are abundant in nature.

Nonmetals u Most nonmetals don’t conduct electricity, are much poorer conductors of heat than metals, and are brittle when solid.

Nonmetals u Many are gases at room temperature; those that are solids lack the luster of metals. u Their melting points tend to be lower than those of metals.

Nonmetals u With the exception of carbon, nonmetals have five, six, seven, or eight valence electrons. u Nonmetals tend to gain electrons in order to achieve the stability of a filled octet.

Nonmetal Reactivity u A nonmetal’s reactivity is its ability to react with another substance. u Reactivity for nonmetals increases as you go left to right and up the periodic table.

Nonmetal Reactivity 2. Consult the “Activity Series of Halogens” in the Chemistry Reference Tables to determine the less active nonmetal. (chlorine) a) fluorine (F 2 ) or chlorine (Cl 2 ) b) chlorine (Cl 2 ) or iodine ( I 2 ) (iodine)

u The metalloids are in pink. Metalloids

Metalloids u Metalloids have some chemical and physical properties of metals and other properties of nonmetals. u In the periodic table, the metalloids lie along the border between metals and nonmetals.

Group Names u Groups 1, 2, and (Group A elements) are called representative (main group) elements.

Group Names u Groups (Group B elements) are called transition elements.

Group 1: The Alkali Metals

u Group 1 elements have one valence electron. u They form 1+ ions after losing the one valence electron.

Group 2: The Alkaline Earth Metals

u Group 2 elements have two valence electrons. u They form 2+ ions after losing the two valence electrons.

Group 17: The Halogens

u Group 17 elements have seven valence electrons. u They form 1- ions after gaining one more electron.

Group 18: The Noble Gases

u Group 18 elements have eight valence electrons, except for helium which only has two. u The noble gases, with a full complement of valence electrons, are generally unreactive.

Valence Electrons

Question How many valence electrons are in an atom of each of the following elements? a) Magnesium (Mg) b) Selenium (Se) c) Tin (Sn) (2) (6) (4)

Question 2. Match each element in Column A with the best matching description in Column B. Each Column A element may match more than one description from Column B.

Question 1. strontium 2. chromium 3. iodine Column A Column B a. halogen b. alkaline earth metal c. representative element d. transition element

Answers 1. strontium 2. chromium 3. iodine b, c d a, c

Periodic Trends

u Because the periodic table relates group and period numbers to valence electrons, it’s useful in predicting atomic structure and, therefore, chemical properties.

Periodic Trends

Atomic Size (Atomic Radius) u The atomic radius of a chemical element is a measure of the size of its atoms, usually the mean or typical distance from the nucleus to the boundary of the surrounding cloud of electrons.

Trends in Atomic Size (Radii) u Atomic size is influenced by two factors. Energy Level – A higher energy level is farther away. Charge on nucleus - More charge (protons) pulls electrons in closer.

Group Trend for Atomic Radii u As you go down a group, each atom has another energy level so the atoms get bigger. H Li Na K Rb

Period Trend for Atomic Radii u As you go across a period, the radius gets smaller. u Atoms are in the same energy level, but as you move across the chart, atoms have a greater nuclear charge (more protons). u Therefore, the outermost electrons are closer.

Period Trend for Atomic Radii NaMgAlSiPSClAr

Question 3. (a) State why atoms get bigger as you go down a group on the periodic table. (b) State why the radius decreases across a period.

Question 4. Choose the element from the pair with the larger atomic radius. a) lithium (Li) or beryllium (Be) (lithium) b) silicon (Si) or tin (Sn) (tin)

Question 5. Choose the element from the pair with the smaller atomic radius. a) silver (Ag) or gold (Au) (silver) b) cesium (Cs) or barium (Ba) (barium)

Ionic Size (Ionic Radius) u Ionic radius is the radius of an atom's ion. u When an atom gains or loses one or more electrons, it becomes an ion.

Ionic Size (Ionic Radius) u Recall that metals tend to lose electrons in order to achieve the stability of a filled octet. u As a result, metals tend to form cations which are positive ions.

Ionic Size (Ionic Radius) u A cation has a smaller radius than its neutral atom.

Ionic Size (Ionic Radius) u Nonmetals tend to gain electrons in order to achieve the stability of a filled octet. u As a result, nonmetals tend to form anions which are negative ions.

Ionic Size (Ionic Radius) u An anion has a larger radius than its neutral atom.

Question 6. Choose the element from the pair with the smaller radius. a) silver (Ag) or the silver ion (Ag 1+ ) (silver ion) b) oxygen (O) or the oxygen ion (O 2- ) (oxygen)

Question 7. For each of the following pairs, predict which atom is larger. a) Mg, Sr b) Sr, Sn c) Ge, Sn d) Ge, Br e) Cr, W (Sr) (Sn) (Ge) (W)

Question (Mg) (S 2- ) (Ba 2+ ) (I-)(I-) (Na + ) 8. For each of the following pairs, predict which atom or ion is larger. a) Mg, Mg 2+ b) S, S 2– c) Ca 2+, Ba 2+ d) Cl –, I – e) Na +, Al 3+

Ionization Energy  Ionization energy ( I E) is the amount of energy required to completely remove an electron from a gaseous atom. u Removing one electron makes a 1+ ion. The energy required to do this is called the first ionization energy.

Ionization Energy

What Determines Ionization Energy (IE) u Greater the nuclear charge (# of protons) means greater IE. u The shorter the distance from the nucleus, the greater the IE.

Ionization Energy u As you go down a group, first IE decreases. u This is because the electron is farther away, thus there is more shielding by the core electrons from the pull of the positive nucleus.

Ionization Energy u All the atoms in the same period have the same energy level. u They have the same shielding, but as you move across the chart there is an increasing nuclear charge because of the increasing number of protons. u Therefore, IE generally increases from left to right.

Question 9. (a) State why ionization energy decreases as you go down a group. (b) State why ionization energy increases across a period.

Question 10. Choose the element from the pair with the greater ionization energy. a) silver (Ag) or iodine ( I ) (iodine) b) oxygen (O) or selenium (Se) (oxygen)

Question 11. Choose the element from the pair with the smaller ionization energy. a) chromium (Cr) or tungsten (W) (tungsten) b) sodium (Na) or magnesium (Mg) (sodium)

Electronegativity u Electronegativity is the tendency for an atom to attract a pair of electrons to itself when it is chemically combined with another element. u Large electronegativity means the atom pulls the electron toward it.

Electronegativity

u Electronegativity decreases down a group. u The farther you go down a group, the farther the electron is away from the nucleus and the more electrons an atom has.

Electronegativity u It is harder to attract extra electrons if the available energy level is far from the nucleus, so electronegativity decreases.

Electronegativity u As you go across a row, electronegativity increases. u Remember the radius of the atoms decreases across the periodic table. u With the smaller size, there is a greater attraction for the nucleus to electrons.

Question 12. (a) State why electronegativity decreases as you go down a group. (b) State why electronegativity increases across a period.

Question 13. Choose the element from the pair with the greater electronegativity. a) sodium (Na) or rubidium (Rb) (sodium) b) selenium (Se) or bromine (Br) (bromine)

Question 14. Choose the element from the pair with the smaller electronegativity. a) magnesium (Mg) or calcium (Ca) (calcium) b) nitrogen (N) or oxygen (O) (nitrogen)

Summary of the Periodic Trends u Moving Left → Right Atomic Radius Decreases Ionization Energy Increases Electronegativity Increases u Moving Top → Bottom Atomic Radius Increases Ionization Energy Decreases Electronegativity Decreases

Electron Affinity Electron affinity is the energy change that accompanies a gaseous atom when it gains an electron to form a gaseous ion.

Metallic Character and Periodic Trends u Recall again that metals tend to lose electrons in order to achieve the stability of a filled octet. u Therefore as metallic character increases toward francium (Fr) ionization energy, electronegativity, and electron affinity decrease.