Periodicity

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Guiding Questions Why is the periodic table so important? Why is the periodic table shaped the way it's shaped? Why do elements combine? Why do elements react? What other patterns are there in the world and how do they help us? Periodic Table Study Questions   1.         Why did chemists make the periodic table? 2.         Why was the table difficult to make? 3.         Why were Dobereiner’s triads of limited use at a periodic table? 4.         What did Newland discover about the elements? 5.         What did Meyer contribute to the development of the periodic table? 6.         What did Mendeleev use as the organizing property for the periodic table? 7.         What problem developed from the use of this property? 8.         What is common to elements in a column of the table? 9.         How did properties change in a row of the table? 10.       What was the significance of gaps in Mendeleev’s periodic table? 11.       What did Moseley use to order the elements in the periodic table? 12.       How did Moseley change the periodic law? 13.       What determines the identity of an element? 14.       Why do elements in a column of the periodic table have similar properties? 15.       With respect to the Periodic Table, what is the meaning of periodic? 16.       What does a row of the Periodic table represent? 17.       What happens to valence electrons as you move left to right in a row? 18.       When determines stability in an atom? 19.       List, from least to most, the stable configurations in an atom. 20.       What determines the column of the periodic table an element is in? 21.       What sublevels are in the outer level of an atom? 22.       What is the maximum number of electrons in the outer level of an atom? 23.       What determines the row and column of the periodic table an element is in? 24.       What are common properties of metals? 25.       What are common properties of non-metals? 26.       What three things can happen to electrons when atoms form compounds? 27.       The configuration of He is 1s2, but it is placed in column 18. Explain this discrepancy. 28.       Hydrogen is obviously not an alkali metal. Why is it in column 1 of the table? 29.       What is necessary for a metalloid to act as a semiconductor?

Periodicity Elements in the PT are arranged in order of increasing atomic number. Elements in the same group - same chemical and physical properties. Across the period - repeating pattern of physical and chemical properties known as periodicity.

Periodic Trends Properties such as Atomic radii and ionic radii First ionisation energy Melting points Electronegativity show periodicity

Atomic Radii (pm) of the Elements Radius : half the distance between neighbouring nuclei (from nucleus to outermost electons) Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Ionic Radius When an atom gains or loses electrons, the radius changes Cations are always smaller than their parent atoms (often losing an energy level) Anions are always larger than their parent atoms (increased e repulsions) Copyright McGraw-Hill 2009

Atomic & ionic radii down a group Atomic radii is determined by 2 opposing factors Shielding effect by the electrons of the inner shell(s) Nuclear charge (due to protons) Moving down the group, both the nuclear charge and shielding effect increase. However, the outer electrons enter new shells. So, although the nucleus gains protons, the electrons are not only further away, but also more effectively screened by an addtional shell of electrons. Shielding effect – makes the atomic radius larger , the result of repulsion between the electrons in the inner shell and those in the outer shell. Nuclear charge – an attractve force that pulls all the electrons closer to the nucleus. With an increase in nuclear charge, atomic radius becomes smaller. Down the group, increase in atomic radius as the nuclear charge increases, due to 2 factors The increase in the no. of complete electron shells between the outer electrons and the nucleus - increasse in the shielding effect of the outer electrons by the inner electrons

Ionic radii for ions of the same charge also increases down a group for the same reason. Atomic radius increases down the group

Copyright McGraw-Hill 2009 Isoelectronic Series Two or more species having the same electron configuration (same number of electrons) but different nuclear charges Size varies significantly Copyright McGraw-Hill 2009

Atomic Radius vs. Atomic Number 0.3 Cs Rb 0.25 K 0.2 Na 3d transition series 4d transition series atomic radius La Li 0.15 Zn Xe Kr 0.1 Cl F 0.05 He H 0 10 20 30 40 50 60 atomic number

Ionic radii across a period The radii of positive ions decrease from Na+ to Al 3+ The radii of positive ions decrease from P3- (phosphide ion) to Cl - The ionic radii increase from the Al 3+ to P3- .

Copyright McGraw-Hill 2009 Explain What do you notice about the atomic radius across a period? Why? What do you notice about the atomic radius down a column? Why? Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 What do you notice about the atomic radius across a period? Why? Atomic radius decreases from left to right across a period due to increasing nuclear charge but no significant increase in the shielding effect. The force of attraction between the negatively charged valence electrons and the positive nucleus increases across the period. What do you notice about the atomic radius down a column? Why? Atomic radius increases down a column of the periodic table because the distance of the electron from the nucleus increases as n increases. Copyright McGraw-Hill 2009

Ionisation Energy (IE) The first ionisation energy is the energy required to remove one electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+. X(g)  X+(g) + e-

Li Li+ e Energy e Li Lithium atom e Li+ e Li Li + e Lithium ion 152 Li 60 Li+ e Energy e 152 Li Lithium atom 152 e 60 Li+ e Li Li + e Lithium ion Lithium atom

Things to remember about IE When talking about first ionization energies (IE), everything must be in gas form IE are measured in kilojoules per mole. All elements have a first ionization energy, even those that do not form cations. What can you conclude if their ionization energy is very high ? It is difficult to lose an electron.

Successive ionisation energies Na(g)  Na+(g) + e Na+(g)  Na2+(g) + e Na2+(g)  Na3+(g) + e

IE of Aluminium What pattern do you notice? What does this suggest about the energy levels?

Multiple Ionization Energies 2745 kJ/mol e- 578 kJ/mol e- 1817 kJ/mol e- Al Al+ Al2+ Al3+ 1st Ionization energy 2nd Ionization energy 3rd Ionization energy • In an atom that possesses more than one electron, the amount of energy needed to remove successive electrons increases steadily. • Define a first ionization energy as (1), a second ionization energy as (2) and in general an nth ionization energy (n) according to the equation E(g)  E+(g) + e-- 1 = 1st ionization energy E+(g)  E2+(g) + e-- 2 = 2nd ionization energy E(n-1)+(g)  En+(g) + e-- n = nth ionization energy The second, third, and fourth ionization energies of aluminum are higher than the first because the inner electrons are more tightly held by the nucleus. Smoot, Price, Smith, Chemistry A Modern Course 1987, page 190

1st IE

Copyright McGraw-Hill 2009 IE1 (kJ/mol) Values for Main Group Elements Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Explain 1. What do you notice about the 1st IE across a period? Why? When moving across the period from left to right, the nuclear charge increases but the shielding effect only increases slightly (since electrons enter the same shell). Hence, the electron shells are pulled progressively closer to the nucleus and it is harder to remove the valence electron. Copyright McGraw-Hill 2009

Explain 2. What do you notice about the 1st IE down a column? Why? The atomic radius increases down the group as additional electrons are added, causing the shielding effect to increase. The further the outer shell is from the nucleus, the smaller the attractive force exerted by the protons in the nucleus. More easily an outer electron can be removed, the lower the ionisation energy.

First ionisation across a period When moving across the period from left to right, the nuclear charge increases but the shielding effect only increases slightly (since electrons enter the same shell). Hence, the electron shells are pulled progressively closer to the nucleus. BUT.........there are exceptions

Be 1s22s2 1st I.E. = 900 kJ mol-1 B 1s22s22px1 1st I.E. = 799 kJ mol-1

Copyright McGraw-Hill 2009 Removal of electrons Copyright McGraw-Hill 2009

Adding of electrons

Discussion B has lower IE than Be same shielding greater nuclear charge

Factors affecting ionisation energy The charge on the nucleus. The distance of the electron from the nucleus. The number of electrons between the outer electrons and the nucleus. Whether the electron is on its own in an orbital or paired with another electron.

Limitation to Bohr’s Model The first IE of the elt 3 (Li) to 10 (Ne) do not increase evenly. There is a need for a more complex model of electron configurations than the Bohr model. Each main energy level is an atom is made up of sub energy levels (subshells).

Plenary - K U I As a result of the lesson today I: Know… Understand… Can use the information in the following other situations….

Electronegativity The electronegativity is the ability of an atom in a covalent bond to attract shared pairs electrons to itself. The greater the electronegativity of an atom, the greater its ability to attract shared pairs of electrons to itself. Electronegativity value is based on the Pauling scale. A value of 4.0 is give to F (most electronegative atom). The least electronegative elements, Ce and Fr both have a value of 0.7

Trends in electronegativity There is an increasing distance between the nucleus and electrons down the group. Hence, the attractive force is decreased. Although the nuclear charge increases down the group, this is counteracted by the increased shielding effect due to additional electron shells.

Trends in melting point Group I Metals are held together by metallic bonding. The strength of metallic bonding decreases because the attractive forces between the delocalised electrons and the nuclues decreases owing to the increase in the distance. The increase in the nuclear charge is counteracted by the increase in shielding. Group 7 As the molecules become large, the attractive forces between them increases with the number of electrons in atoms or molecules. Metals are composed of a lattice of positive ions surrounded by delocalised electrons which move between the ions.

The melting point decreases down group 1 The melting point increases down group 7

Periodic Trends in Chemical Properties of Main Group Elements IE and EA enable us to understand types of reactions that elements undergo and the types of compounds formed Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 General Trends in Chemical Properties Elements in same group have same valence electron configuration; similar properties Same group comparison most valid if elements have same metallic or nonmetallic character Group 1A and 2A; Group 7A and 8A Careful with Group 3A - 6A Copyright McGraw-Hill 2009

Chemical Properties Group I alkali metals Li, Na and K contain 1 valence electron. Reactive metals, stored under liquid paraffin to prevent them from reacting with air. Readily lose their valence electron -good reducing agent Reactivity increases down the group Soft metals of low density with a low melting point. Form M+ cations. Relatively low 1st IE and are therefore chemically reactive. Strong reducing agent and ther ions are hard to reduce. Reactivity increases down the group and correlates with a decrease in the 1st IE, due to the increasing distance between the nucleus and the valence electron. Atomic and ioni radii increase and eletronegativity and melting point decreases down the group due to the presence of additional eletron shell. Oxygen with heated metal 2M(s) + ½ O2(g)  M2O(s) Halogen with heated group 1 metal M(s) + ½ X2(g)  MX(s) Water with metal M(s) + H2O(l)  MOH(aq)+ ½ H2(g)

React with water to form an alkali solution of the metal hydroxide and hydrogen gas. 2Li(s) + 2H2O(l)  LiOH(aq) + H2(g) Lithium floats and reacts quietly (ii) 2Na(s) + 2H2O(l)  NaOH(aq) + H2(g) Sodium melts into a ball which darts around on the surface 2Ks) + 2H2O(l)  KOH(aq) + H2(g) Heat generated from the reaction with potassium ignites the hydrogen.

React readily with chlorine, bromine and iodine to form ionic salts, e React readily with chlorine, bromine and iodine to form ionic salts, e.g. 2Na(s) + Cl2(g)  2NaCl(s) 2K(s) + Br2(l)  2KBr(s) 2Ks) + I2(g)  2LiI(s)

Chemical Properties Chlorine is a stronger oxidizing agent than bromine, so can remove the electron from bromide ions in solution to form chloride ions and bromine. Similarly, both chlorine and bromine can oxidize iodide ions to form iodine. Cl2(aq) + 2Br-(aq)  2Cl-(aq) + Br2(aq) Cl2(aq) + 2I-(aq)  2Cl-(aq) + I2(aq) Br2(aq) + 2I-(aq)  2Br-(aq) + I2(aq) Halogens are a group of reactive non-metals n group 7. They all form X- ions Reactivity increases up the group. This correlates with an increase in the 1st eletron affnity due to decreasing dstance between the nucleus and the inoming elecron. Displacement reactions X2(aq) + 2Y- (aq)  2X- (aq) + Y2(aq) X represents a more reactive halogen (more powerful oxidising agent) than Y Reaction with water X2(aq) + H2O(l) HOX (aq) + H+(aq) + X-(aq) Reaction with group 1 metals ½ X2(g) + M(s)  MX(s) Precipitation reactions X- (aq) + Ag+ (aq)  AgX(s)

Test for halide ions The presence of halide ions in solution can be detected by adding silver nitrate solution. Ag+(aq ) + X- (aq)  AgX(s) X = Cl,Br or I light AgCl white Ag(s) + ½ X2 AgBr cream AgI yellow