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

Isoelectronic Species Isoelectronic species are atoms and ions that have the same number of electrons The large increase in size from Al3+ to P3- is due to the presence of additional electron shell which causes a large increase in the shielding effect resullting in an increase in the ionic radius. Species Na+ Mg2+ Al3+ Nuclear charge +11 +12 +13 Number of electrons 10 Ionic configuration 98 65 45 Species P3- S2- Cl- Nuclear charge +15 +16 +17 Number of electrons 18 Ionic configuration 212 190 181

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

Why is the 2nd IE always higher than the 1st IE? Positive Ion When one electron is removed, a positive ion is formed. It attracts the negatively charged electron more strongly than the neutral atom. Hence more energy is needed to remove the electron Repulsion between electrons When one electron is remove, there is less repulsion between the remaining electrons. Hence, they are pulled closer to the nucleus. more energy is needed to remove the electron

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 Exception in variation of IE across the period ( 1st IE of B is lower than that for Be ) B Be The electron to be removed from B is from 2p sub level which is at a higher level than 2s The energy required to remove the electron from 2p in B is lesser than that required to remove the electron from 2s in Be. Copyright McGraw-Hill 2009

Exception in variation of IE across the period ( 1st IE of O is lower than that for N ) O has 2 electrons paired up in the same p orbital, unlike N where each electron is singly occupied in each p orbital. It is easier to remove the electron in the same orbital because of the repulsion from the other electron.

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….

Atoms in a covalent bond are held together by electrostatic forces of attraction between positively charged nuclei and negatively charged shared electrons.

Pure covalent bond The electronegativity is the ability of an atom in a covalent bond to attract shared pairs electrons to itself. When two atoms bonded by a covalent bond have the same electronegativity the electrons will be equally shared. This is pure (non-polar) covalent bonding. For example, the chlorine molecule, Cl2. Pure covalent bond is only found in elements

Polar covalent bond Covalent bonds with unequal electron sharing are called polar covalent bonds. The atom with the higher electronegativity will have a greater share of the electrons. result in the atom having a slight negative charge compared to the other atom which will have a slight positive charge. In most covalent compound, the bonding is polar covalent http://www.bbc.co.uk/bitesize/higher/chemistry/energy/bsp/revision/1/

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 given to F (most electronegative atom). The least electronegative elements, Cs and Fr both have a value of 0.7 Electronegativity is a relative value, not absolute value. Not a precise value that each atom has. The difference in electronegativity between 2 bonding atoms.

Electronegativity down the group Do noble gases have any electronegative value? There is an increasing distance between the nucleus and electrons down the group. Hence, the attractive force is decreased. [compare HF and HCl] Although the nuclear charge increases down the group, this is counteracted by the increased shielding effect due to additional electron shells.

The electronegativity decreases down the group Group 1 & 2 metals are sometimes called electropositive elements Group 5,6 & 7 non-metals are sometimes called electronegative elements The electronegativity decreases down the group

Electronegativity across a period Higher nuclear charge across the period with no significant change in shielding when the number of shells is the same. [compare HN and HF]

The electronegativity increases across the period Group 1 & 2 metals are sometimes called electropositive elements Group 5,6 & 7 non-metals are sometimes called electronegative elements The electronegativity increases across the period

Melting Points Depends on the structure of the element and the type of attractive forces holding the atoms together.

Trends in melting point Group I (Metallic bonding) Metals are held together by electrostatic attraction between the positive ions in the lattice and the delocalised electrons. Metals are composed of a lattice of positive ions surrounded by delocalised electrons which move between the ions.

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. Less energy is required to break apart the lattice

Group 7 As the molecules become large, the van der Waals’ attraction between the diatomic molecules increases. More energy is required to separate the molecules from each other

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

Variation in melting point across period 3

Metallic Bonding

Melting point for sodium, magnesium and aluminium all metals have metallic bonding, in which positive metal ions are attracted to delocalised electrons. From sodium to aluminium: the charge on the metal ions increases from +1 to +3 (with magnesium at +2) the number of delocalised electrons increases so the strength of the metallic bonding increases and the melting points and boiling points increase. http://www.chemguide.co.uk/atoms/bonding/metallic.html

Covalent Bonding – tetrahedral structure

Silicon a metalloid . has giant covalent bonding silicon atom is tetrahedrally bonded to 4 other silicons by a covalent bond. very high melting point and boiling point because: all the silicon atoms are held together by strong covalent bonds which need a very large amount of energy to be broken.

Van der Waals – simple molecules

Phosphorus, sulphur, chlorine and argon all non-metals, and they exist as small, separate molecules. Phosphorus, sulphur and chlorine exist as simple molecules, with weak van der Waals’ forces of attraction between the molecules. These are instantaneous attractions between the electrons and nucleus’ of different atoms. Argon exists as separate atoms (it is monatomic).

Their melting and boiling points are very low because: when these four substances melt or boil, it is the van der Waals’ forces between the molecules which are broken which are very weak bonds so little energy is needed to overcome them.

phosphorus exists as P4 molecules sulphur exists as S8 molecules Sulphur has a higher melting point and boiling point than the other three because: phosphorus exists as P4 molecules sulphur exists as S8 molecules chlorine exists as Cl2 molecules argon exists individual Ar atoms the strength of the van der Waals’ forces decreases as the size of the molecule decreases so the melting points and boiling points decrease in the order S8 > P4 > Cl2 > Ar The more atoms in a molecule the stronger the Van der Waals attraction – as there are more electrons in the molecule. S8, P4, Cl2

Complete this table using arrows! Across a Period Down a Group Nuclear Charge/ No. of protons Shielding Effect Distance from nucleus Ionisation Energy

Which element in the periodic table would you expect to have the BIGGEST ionisation energy SMALLEST ionisation energy Why?

Draw a Sodium atom and a Sodium ion. Which is going to be smaller and WHY (3 marks).

Match name to picture to description A: Covalent Bonding X : Instantaneous dipole attraction – very weak 1 Y: Strong attraction between ion and delocalised electrons B: Van der Waals Forces 2 Z : Strong bond between atoms by sharing electrons 3 C: Metallic Bonding

Why does aluminium have a higher melting point than sodium? Why does silicon have the highest melting point? Why does Sulphur have a higher melting point than Chlorine?

Now complete past paper questions….

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 1 and 2; Group 7 and 8 Careful with Group 3 - 6 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)

Reaction with oxygen React with oxygen to form metal oxides (i) 4Li(s) + O2(g)  2Li2O(s) (ii) 4Na(s) + O2(g)  2Na2O(s)

Reaction with water 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.

Reaction with halogens 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. Displacement Reactions 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)

Oxides of Period 3

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

Summary of Periodic Trends