TOPIC 3 PERIODICITY 3.2 PERIODIC TRENDS
Elements show trends in their physical and chemical properties across periods and down groups. NATURE OF SCIENCE (3.1) Looking for patterns – the position of an element in the periodic table allows scientists to make accurate predictions of its physical and chemical properties. This gives scientists the ability to synthesize new substances based on the expected reactivity of elements. ESSENTIAL IDEA
Industrialization has led to the production of many products that cause global problems when released into the environment. INTERNATIONAL-MINDEDNESS
The predictive power of Mendeleev’s Periodic Table illustrates the “risk-taking” nature of science. What is the demarcation between scientific and pseudoscientific claims? The Periodic Table is an excellent example of classification in science. How does classification and categorization help and hinder the pursuit of knowledge? THEORY OF KNOWLEDGE
When elements are arranged according to atomic number, there is a periodic pattern in their properties. • Periodic means the pattern is repeated over and over. • Remember that electrons play the biggest part in determining the physical and chemical properties of an element. • The periodicity of the elements is reflected in their physical properties. PERIODICITY
Vertical and horizontal trends in the periodic table exist for atomic radii, ionic radius, ionization energy, electron affinity and electronegativity. UNDERSTANDING/KEY IDEA3.2.A
You only have to know examples of general trends across periods and down groups. For ionization energy, the discontinuities in the increase across a period should be known. (These are the two exceptions discussed in Chapter 2.) GUIDANCE
The concept of effective nuclear charge helps explain the trends in chemical and physical properties. • The “nuclear charge” is given by the number of protons and it increases by one between successive elements as a proton is added to the nucleus. • The outer shell electrons which determine many of the chemical and physical properties do not experience the full attraction of the nucleus because they are shielded from the nucleus and repelled by the inner shell electrons. EFFECTIVE NUCLEAR CHARGE
The “effective charge” experienced by the outer shell electrons is less than the full nuclear charge. • As you “go down” a group in the Periodic Table, the effective nuclear charge remains about the same because the inner shell electrons “shield” the outer electrons from the nucleus. • Sodium has 11 electrons with 1 e- in the outer shell. The outer shell electron is shielded from the 11 protons by the 10 interior electrons in the first and second energy levels. • As you go from left to right, the effective nuclear charge increases because as you are adding protons, you are still in the same energy level so the inner electrons are not increasing. • To summarize, the effective nuclear charge increases across a period, but remains about the same down a group.
The “atomic radius” is not “how big the atom is” but is measured as half the distance between neighboring nuclei. • The atom does not have sharp boundaries so it is difficult to measure its size. • The atomic radius increases as you go down a Group because you are adding energy levels. • The atomic radius decreases across a Period because the effective nuclear charge increases which increases the attraction between the nucleus and the outer shell electrons. ATOMIC RADIUS
1. Positive ions (cations) are smaller than their parent atoms because they have lost one or more electrons. • 2. Negative ions (anions) are larger than their parent atoms because they have gained one or more electrons. Increasing electron repulsions also cause the radius to increase. • 3. The ionic radius decreases from Groups 1 to 4 for the positive ions because of increasing effective nuclear charge. The increased attraction between the nucleus and the electrons pulls the outer shell electrons closer to the nucleus. IONIC RADIUS
4. The ionic radius decreases from Groups 14 to 17 for the negative ions for the same reason as mentioned in trend 3. There is a jump however in size when you move from positive to negative ions within a period. • 5. The ionic radii increases as you move down a Group due to increasing energy levels.
The first ionization energy of an element is the energy required to remove one mole of electrons from one mole of gaseous atoms. • Electronegativity is the measure of the tendency of an atom in a molecule to attract a shared pair of electrons towards itself. • Electron affinity is the energy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous ions.
First ionization energies are a measure of the attraction between the nucleus and the outer shell electron. • 1. Ionization energies increase across a period due to increasing nuclear charge which makes it more difficult to remove the outer shell electron. • 2. Ionization energies decrease down a group. This is due to the distance from the nucleus and the shielding by the inner shell electrons. These 2 factors make it easier to remove the outer shell electrons the farther they are away from the nucleus. IONIZATION ENERGIES
1. Remember from Chapter 2 that the Group 3 elements have lower energies than Group 2 elements because of electron configuration. • The Group 3 elements have the general configuration of ns2, np1 and Group 2 elements are ns2. It is easier to remove the “p” electron than the “s” electrons. • 2. Group 16 elements have lower energies than Group 15 elements also due to electron configuration. • The Group 16 elements have a doubly occupied “p” orbital, unlike the Group 15 elements which have three singly occupied “p”orbitals. • It is easier to pull an electron from a doubly occupied orbital because of the added electron repulsion. IONIZATION ENERGY EXCEPTIONS
The electronegativity of an element is a measure of the ability of its atoms to attract electrons in a covalent bond. • It is also a measure of the attraction between the nucleus and the outer shell bonding electrons. • An element with a high electronegativity has strong electron pulling power. • An element with a low electronegativity has weak electron pulling power. • Fluorine is the most electronegative element and is located in the top right corner. • The least electronegative element is at the bottom left. ELECTRONEGATIVITY
Electronegativity follows the same trend as ionization energy. • 1. It increases as you move from left to right due to increasing effective nuclear charge. • 2. It decreases as you go down a group because the outer shell electrons are farther away from the nucleus which decreases the attraction. • Even though the trends are the same, ionization energy and electronegativity are distinct properties. • Ionization energies can be measured directly and are a property of gaseous atoms. • Electronegativity is a property of an atom in a molecule and its values are derived indirectly from experimental bond energy data. ELECTRONEGATIVITY
The elements with the highest electronegativities are located in the top right. • The elements with the lowest electronegativities are located in the bottom left.
Trends in electron affinity are not as clear as for the other trends studied. • 1. Electron affinity is an exothermic process so the values are negative. • 2. Electron affinity values generally get more negative as you go from left to right in a period which means that the halogens really want to add that extra electron. • 3. The patterns vary down a group and do not show the same clear trends as do atomic radii, ionization energy and electronegativity. ELECTRON AFFINITY
You should know the trends for the following group reactions: alkali metals with water, alkali metals with halogens and halogens with halide ions. GUIDANCE
Alkali metals react with water to form hydrogen gas and the metal hydroxide. • They are called alkali metals because they create a basic or alkaline solution when reacted with water. • The reaction with water becomes more vigorous as you descend a group because the ionization energy decreases and positive ions are formed more readily. • Li reacts slowly, releases H2 but keeps is shape and floats. • Na reacts with a vigorous release of H2 which releases enough heat to melt the remaining Na that floats on the surface. • K reacts vigorously to produce enough heat to ignite the H2 produced and creates a lilac flame which dances on the surface. GROUP 1 – REACTIVITY WITH WATER
Halogens react with Group 1 alkali metals to form ionic or metallic halides. • The metal loses its electron to the halide to form a halide ion X- and a positive ion M+. • After the electron is transferred, the ions are pulled together by the strong attraction between the positive and negative charges. • The most vigorous reactions occur between elements that are farthest apart on the periodic table. (Bottom left with top right.) HALOGENS REACT WITH GROUP 1 METALS
The more reactive halogen displaces the ions of the less reactive halogens. F>Cl>Br>I • This is seen in single replacement reactions. • The smaller the halogen, the more reactive it is. • Color changes indicate displacement reactions. • When Br2 is displaced by chlorine, the solution turns from clear to orange. • When I2 is displaced by bromine, the color darkens even more. It turns purple when mixed with a hydrocarbon. • Halogens also form insoluble salts with silver which can help identify the halide. DISPLACEMENT REACTIONS OF HALOGENS
Trends in metallic and non-metallic behavior are due to the trends stated in 3.2.A. UNDERSTANDING/KEY IDEA3.2.B
Be able to predict and explain the metallic and non-metallic behavior of an element based on its position in the periodic table. APPLICATION/SKILLS
Metallic character increases down a group and decreases from left to right or across a period. • Metals tend to have low ionization energies as they tend to lose electrons in a chemical reaction. • Non-metals have high electron affinities as they tend to gain electrons in chemical reactions.
The chemical properties of an element are largely determined by the number of electrons in the outer shell.
The reactivity of elements in other groups is explained by their unstable incomplete electron shells. • They lose or gain electrons to achieve the electron arrangement of the nearest noble gas. • Elements in Groups 1 to 3 lose electrons to adopt the arrangement of the nearest noble gas with a lower atomic number. • They are usually metals. • Elements in Groups 15 to 17 gain electrons to adopt the electron arrangement of the nearest noble gas to their right. • They are usually non-metals. • The metalloids show intermediate properties.
Be able to discuss similarities and differences in the properties of elements in the same group, with reference to alkali metals (group 1) and halogens (group 17). APPLICATION/SKILLS
The alkali metals are silvery metals that are too reactive to be found in nature – they are stored in oil to prevent contact with air and water. • The first 3 elements have the following properties: • Physical • Good conductors of electricity • Low densities • Have grey shiny surfaces when cut with a knife • Chemical • They are very reactive metals. • They form ionic compounds with non-metals. GROUP 1 – ALKALI METALS
They form single charged ions with a 1 plus charge. (M+) • Reactivity increases down a group because ionization energy decreases and the outer shell electron is more easily removed. • Their ability to conduct electricity is due to the mobility of the outer shell electron.
The halogens exist as diatomic elements. • Physical properties: • They are colored. • They show a gradual change from gases (F2 and Cl2) to liquid (Br2) to solids (I2 and At2). • Chemical properties: • They are very reactive non-metals. • Reactivity decreases down the group. • They form ionic compounds with metals or covalent compounds with other non metals. • Reactivity decreases down the group as the atomic radii increases and the attraction for outer electrons decreases. GROUP 17 – HALOGENS
Group 18 – the Noble Gases – contains the least reactive elements. • With the exception of Helium, they have complete outer shells of 8 electrons: a stable octet. • Their lack of reactivity is explained by their inability to lose or gain electrons. • Noble gases – are colorless gases, exist as single atoms (monatomic), and are very unreactive. GROUP 18 – NOBLE GASES
Melting points depend on both the type of bonding and the structure. • Melting points decrease down Group 1 because the metallic structures are held together by the attractive forces of delocalized outer shell electrons and the positive ions. This attraction decreases with distance. • Melting points increase down Group 17 because the molecular structures are held together by van der Waal’s forces which increase with the number of electrons in a molecule. • Melting points rise across a period until they reach Group 14, then they decrease to reach a minimum at Group 18. • All period 3 elements are solids at room temperature except for chlorine and argon. MELTING POINTS
Oxides change from basic through amphoteric to acidic across a period. UNDERSTANDING/KEY IDEA3.2.C
When elements bond with oxygen, oxides are formed. • Ionic compounds are formed between a metal and a nonmetal so the oxides of Na, Mg and Al have giant ionic structures. (high melting pt) • Covalent compounds are formed between non-metals so the oxides of P, S and Cl are molecular covalent. (low melting pt) • The oxide of the metalloid silicon has a giant covalent structure. (high melting pt) BONDING OF PERIOD 3 OXIDES
The ionic character of a compound depends upon the difference in electronegativity between its elements. • Oxygen has an electronegativity of 3.5 so the ionic character of the oxides decreases from left to right as the electronegativity values of the Period 3 elements approach this value. • The oxides become more ionic down a group as electronegativity decreases. • Electrical conductivity is a measure of ionic character and can only be measured in the molten state when electrons are free to move.
Metallic elements which form ionic oxides are basic. • Non-metallic elements which form covalent or non-metallic oxides are acidic. • Na and Mg oxides are basic. • Al and Si oxides are amphoteric – show both acidic and basic properties. • P, S, and Cl oxides are acidic. ACID-BASE CHARACTER OF PERIOD 3 OXIDES
Be able to construct equations to explain the pH changes for reactions of Na2O, MgO, P4O10 and the oxides of nitrogen and sulfur with water. APPLICATION/SKILLS
Basic oxides dissolve in water to form a basic or alkaline solution due to the presence of hydroxide ions. These reactions produce solutions with pH’s above 7 (basic). Na2O + H2O → 2NaOH MgO + H2O → Mg(OH)2 BASIC OXIDES
The non-metallic oxides react readily with water to produce acidic solutions which have pH’s below 7. • P4O10 + 6H2O → 4H3PO4 • P4O6 + 6H2O → 4H3PO3 • SO3 + H2O → H2SO4 • SO2 + H2O → H2SO3 • 2NO2 + H2O → HNO2 + HNO3 ACIDIC OXIDES
Aluminum oxide does not affect the pH when added to water because it is basically insoluble. • It shows amphoteric properties because it can act like an acid or a base. • It behaves as a base when it reacts with sulfuric acid. • It behaves as an acid when it reacts with bases. AMPHOTERIC OXIDES
International Baccalaureate Organization. Chemistry Guide, First assessment 2016. Updated 2015. Brown, Catrin, and Mike Ford. Higher Level Chemistry. 2nd ed. N.p.: Pearson Baccalaureate, 2014. Print. Most of the information found in this power point comes directly from this textbook. The power point has been made to directly complement the Higher Level Chemistry textbook by Catrin and Brown and is used for direct instructional purposes only. Citations