Atomic Structure & the periodic table Starter: Draw the electron arrangement for carbon, nitrogen, lithium, oxygen, helium.

Ionisation energies Key Words: Subshells Orbitals Principle quantum number Ionisation energy Ionisation Objectives: Discuss Ionisation energy Outcomes: D: recall and understand the definition of ionisation energies of gaseous atoms A-C: Understand that they are endothermic processes

Energy levels & electron shells Electrons in an atom are arranged in a series of shells Shells: Each shell ins described by a principle quantum number:

The larger the value of n, the further from the nucleus you are likely to find the electron:

Ionisation Energies: Ionisation: the complete removal of an electron from an atom It is an endothermic process Since energy is needed to overcome the attractive force between the electron and the nucleus. Ionisation energy: the amount of energy needed to remove an electron from its atom can be measures by increasing voltage applied to a gas until it conducts electricity & emits light – which tells you an electron has been freed

Ground state: the lowest energy state for an atom The energy needed to remove the one electron from Hydrogen in its ground state is normally quoted for 1 mole and is the ionisation energy of hydrogen First ionisation energy: energy needed to remove the first electron from an atom Second ionisation energy: energy needs to remove the second electron from an atom

The first ionisation is a measure of how tightly or loosely an outer electron is attracted to the positive nucleus. The more easily an electron is removed, the more reactive an atom will be. Total energy required to remove electrons  add the 1st & 2nd ionisation energies together. The different energies needs to remove the 1st & subsequent electrons confirm that electrons are found on different energy levels

Maximum electrons per shell THE BOHR ATOM Ideas about the structure of the atom have changed over the years. The Bohr theory thought of it as a small nucleus of protons and neutrons surrounded by circulating electrons. Each shell or energy level could hold a maximum number of electrons. The energy of levels became greater as they got further from the nucleus and electrons filled energy levels in order. The theory couldn’t explain certain aspects of chemistry. Maximum electrons per shell 1st shell 2 2nd shell 8 3rd shell 18 4th shell 32 5th shell 50

Subshells: Quantum mechanics also tells us that each shell may contain subshells. Subshells: regions of differing energy within a shell, shown by letters: s, p, d, f, g. The following subshells are available in each shell:

Shell 1 is closest to the nucleus so it takes the most energy to remove electrons from this shell Electrons in the lowest energy subshells are closest to the nucleus s (lowest energy) < p < d Each type of subshell contains one or more orbitals Orbitals: the region where the electrons are most likely to be found. They hold a maxiumum of 2 electrons

All orbitals in a particular subshell are at the same energy level As n increases, the energy gap between successive shells gets smaller Due to this, orbitals in neighbouring shells may overlap The 3d orbital has an energy level above that of 4s orbital but below 4p orbital

PRINCIPAL ENERGY LEVELS INCREASING ENERGY / DISTANCE FROM NUCLEUS LEVELS AND SUB-LEVELS PRINCIPAL ENERGY LEVELS During studies of the spectrum of hydrogen it was shown that the energy levels were not equally spaced. The energy gap between successive levels got increasingly smaller as the levels got further from the nucleus. The importance of this is discussed later. 4 3 INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 1

PRINCIPAL ENERGY LEVELS INCREASING ENERGY / DISTANCE FROM NUCLEUS LEVELS AND SUB-LEVELS PRINCIPAL ENERGY LEVELS SUB LEVELS During studies of the spectrum of hydrogen it was shown that the energy levels were not equally spaced. The energy gap between successive levels got increasingly smaller as the levels got further from the nucleus. The importance of this is discussed later. A study of Ionisation Energies and the periodic properties of elements suggested that the main energy levels were split into sub levels. Level 1 was split into 1 sub level Level 2 was split into 2 sub levels Level 3 was split into 3 sub levels Level 4 was split into 4 sub levels 4 3 INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 1 CONTENTS

ORBITALS An orbital is... a region in space where one is likely to find an electron. Orbitals can hold up to two electrons as long as they have opposite spin; this is known as PAULI’S EXCLUSION PRINCIPAL. Orbitals have different shapes...

ORBITALS An orbital is... a region in space where one is likely to find an electron. Orbitals can hold up to two electrons as long as they have opposite spin; this is known as PAULI’S EXCLUSION PRINCIPAL. Orbitals have different shapes... ORBITAL SHAPE OCCURRENCE s spherical one in every principal level p dumb-bell three in levels from 2 upwards d various five in levels from 3 upwards f various seven in levels from 4 upwards

DO NOT CONFUSE AN ORBITAL WITH AN ORBIT ORBITALS An orbital is... a region in space where one is likely to find an electron. Orbitals can hold up to two electrons as long as they have opposite spin; this is known as PAULI’S EXCLUSION PRINCIPAL. Orbitals have different shapes... ORBITAL SHAPE OCCURRENCE s spherical one in every principal level p dumb-bell three in levels from 2 upwards d various five in levels from 3 upwards f various seven in levels from 4 upwards An orbital is a 3-dimensional statistical shape showing where one is most likely to find an electron. Because, according to Heisenberg, you cannot say exactly where an electron is you are only able to say where it might be found. DO NOT CONFUSE AN ORBITAL WITH AN ORBIT

s orbitals SHAPES OF ORBITALS spherical one occurs in every principal energy level

p orbitals SHAPES OF ORBITALS dumb-bell shaped three occur in energy levels except the first

d orbitals SHAPES OF ORBITALS various shapes five occur in energy levels except the first and second

ORDER OF FILLING ORBITALS INCREASING ENERGY / DISTANCE FROM NUCLEUS 1 1s 2 2s 2p 4s 3 3s 3p 3d 4 4p 4d 4f PRINCIPAL ENERGY LEVELS SUB LEVELS Orbitals are not filled in numerical order because the principal energy levels get closer together as you get further from the nucleus. This results in overlap of sub levels. The first example occurs when the 4s orbital is filled before the 3d orbitals.

ORDER OF FILLING ORBITALS INCREASING ENERGY / DISTANCE FROM NUCLEUS 1 1s 2 2s 2p 4s 3 3s 3p 3d 4 4p 4d 4f PRINCIPAL ENERGY LEVELS SUB LEVELS 1 1s 2 2s 2p 3d 3 3s 3p 4s 4 4p 4d 4f PRINCIPAL ENERGY LEVELS SUB LEVELS Orbitals are not filled in numerical order because the principal energy levels get closer together as you get further from the nucleus. This results in overlap of sub levels. The first example occurs when the 4s orbital is filled before the 3d orbitals.

Practice:

Orbitals & shells Key Words: Electron spin Electronic configuration Objectives: Electrons spins & filling orbitals Outcomes: D: recall electrons populate orbits singly before pairing up A-C: Understand electron spin - predict the electronic structure & configuration of atoms of hydrogen to krypton inclusive using 1s …notation and electron-in-boxes notation

Electron Spin: Electron Spin: the rotation of electrons clockwise or anticlockwise creating a magnetic field An electron can spin either clockwise or anticlockwise – and because it is moving, it creates a magnetic field This can be represented by using a small arrow ( ) or ( ) – so showing spins in opposite directions

2 electrons in the same orbital cannot have the same spin This means each orbital can have a maximum of 2 electrons, having opposite spins

How would you draw the boxes? Filling the orbitals: Electronic configuration: the arrangement of electrons in an atom in their subshells and orbitals E.g: Hydrogen in it ground state has one electron  1s1 Therefore: Helium: 1s2 Lithium: 1s22s1 Be: 1s22s2 How would you draw the boxes?

Note: the empty p orbitals are shown. It doesn’t matter which 2p orbital is filled first as they all have the same energy

Hund’s Rule

Writing electronic configurations: For an ion: you simply add or subtract the right number of electrons from the outer shell Remember Hund’s Rule when removing electrons: one electron comes out of each completely filled orbital in the outer shell before any unpaired electrons and removed. E.g: O2- is: 1s22s22p6

THE ‘AUFBAU’ PRINCIPAL 4 4p 4d 4f This states that… “ELECTRONS ENTER THE LOWEST AVAILABLE ENERGY LEVEL” 3 3s 3p 3d The following sequence will show the ‘building up’ of the electronic structures of the first 36 elements in the periodic table. Electrons are shown as half headed arrows and can spin in one of two directions or s orbitals p orbitals d orbitals 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p 1 1s

1s1 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f HYDROGEN 4p 4d 4f HYDROGEN 1s1 Hydrogen atoms have one electron. This goes into a vacant orbital in the lowest available energy level. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS ‘Aufbau’ Principle 2 2s 2p 1 1s

1s2 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f HELIUM 1s2 Every orbital can contain 2 electrons, provided the electrons are spinning in opposite directions. This is based on... PAULI’S EXCLUSION PRINCIPLE The two electrons in a helium atom can both go in the 1s orbital. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p 1 1s ‘Aufbau’ Principle

1s2 2s1 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f 4p 4d 4f LITHIUM 1s2 2s1 1s orbitals can hold a maximum of two electrons so the third electron in a lithium atom must go into the next available orbital of higher energy. This will be further from the nucleus in the second principal energy level. The second principal level has two types of orbital (s and p). An s orbital is lower in energy than a p. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p 1 1s ‘Aufbau’ Principle

1s2 2s2 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f 4p 4d 4f BERYLLIUM 1s2 2s2 Beryllium atoms have four electrons so the fourth electron pairs up in the 2s orbital. The 2s sub level is now full. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p ‘Aufbau’ Principle 1 1s

1s2 2s2 2p1 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f 4d 4f BORON 1s2 2s2 2p1 As the 2s sub level is now full, the fifth electron goes into one of the three p orbitals in the 2p sub level. The 2p orbitals are slightly higher in energy than the 2s orbital. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p ‘Aufbau’ Principle 1 1s

1s2 2s2 2p2 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f 4d 4f CARBON 1s2 2s2 2p2 The next electron in doesn’t pair up with the one already there. This would give rise to repulsion between the similarly charged species. Instead, it goes into another p orbital which means less repulsion, lower energy and more stability. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p HUND’S RULE OF MAXIMUM MULTIPLICITY 1 1s

1s2 2s2 2p3 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f 4d 4f NITROGEN 1s2 2s2 2p3 Following Hund’s Rule, the next electron will not pair up so goes into a vacant p orbital. All three electrons are now unpaired. This gives less repulsion, lower energy and therefore more stability. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p HUND’S RULE OF MAXIMUM MULTIPLICITY 1 1s

1s2 2s2 2p4 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f 4d 4f OXYGEN 1s2 2s2 2p4 With all three orbitals half-filled, the eighth electron in an oxygen atom must now pair up with one of the electrons already there. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS ‘Aufbau’ Principle 2 2s 2p 1 1s

1s2 2s2 2p5 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f 4d 4f FLUORINE 1s2 2s2 2p5 The electrons continue to pair up with those in the half-filled orbitals. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p 1 1s

1s2 2s2 2p6 THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4f 4d 4f NEON 1s2 2s2 2p6 The electrons continue to pair up with those in the half-filled orbitals. The 2p orbitals are now completely filled and so is the second principal energy level. In the older system of describing electronic configurations, this would have been written as 2,8. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p 1 1s

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f SODIUM - ARGON With the second principal energy level full, the next electrons must go into the next highest level. The third principal energy level contains three types of orbital; s, p and d. The 3s and 3p orbitals are filled in exactly the same way as those in the 2s and 2p sub levels. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p ‘Aufbau’ Principle 1 1s

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f SODIUM - ARGON Na 1s2 2s2 2p6 3s1 Mg 1s2 2s2 2p6 3s2 Al 1s2 2s2 2p6 3s2 3p1 Si 1s2 2s2 2p6 3s2 3p2 P 1s2 2s2 2p6 3s2 3p3 S 1s2 2s2 2p6 3s2 3p4 Cl 1s2 2s2 2p6 3s2 3p5 Ar 1s2 2s2 2p6 3s2 3p6 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p Remember that the 3p configurations follow Hund’s Rule with the electrons remaining unpaired to give more stability. 1 1s

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f POTASSIUM 1s2 2s2 2p6 3s2 3p6 4s1 In numerical terms one would expect the 3d orbitals to be filled next. However, because the principal energy levels get closer together as you go further from the nucleus coupled with the splitting into sub energy levels, the 4s orbital is of a LOWER ENERGY than the 3d orbitals so gets filled first. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p 1 1s ‘Aufbau’ Principle

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f CALCIUM 1s2 2s2 2p6 3s2 3p6 4s2 As expected, the next electron pairs up to complete a filled 4s orbital. This explanation, using sub levels fits in with the position of potassium and calcium in the Periodic Table. All elements with an -s1 electronic configuration are in Group I and all with an -s2 configuration are in Group II. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p 1 1s ‘Aufbau’ Principle

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f SCANDIUM 1s2 2s2 2p6 3s2 3p6 4s2 3d1 With the lower energy 4s orbital filled, the next electrons can now fill the 3d orbitals. There are five d orbitals. They are filled according to Hund’s Rule - BUT WATCH OUT FOR TWO SPECIAL CASES. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p HUND’S RULE OF MAXIMUM MULTIPLICITY 1 1s

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f TITANIUM 1s2 2s2 2p6 3s2 3p6 4s2 3d2 The 3d orbitals are filled according to Hund’s rule so the next electron doesn’t pair up but goes into an empty orbital in the same sub level. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS HUND’S RULE OF MAXIMUM MULTIPLICITY 2 2s 2p 1 1s

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f VANADIUM 1s2 2s2 2p6 3s2 3p6 4s2 3d3 The 3d orbitals are filled according to Hund’s rule so the next electron doesn’t pair up but goes into an empty orbital in the same sub level. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS HUND’S RULE OF MAXIMUM MULTIPLICITY 2 2s 2p 1 1s

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f CHROMIUM 1s2 2s2 2p6 3s2 3p6 4s1 3d5 One would expect the configuration of chromium atoms to end in 4s2 3d4. To achieve a more stable arrangement of lower energy, one of the 4s electrons is promoted into the 3d to give six unpaired electrons with lower repulsion. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p HUND’S RULE OF MAXIMUM MULTIPLICITY 1 1s

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f MANGANESE 1s2 2s2 2p6 3s2 3p6 4s2 3d5 The new electron goes into the 4s to restore its filled state. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS HUND’S RULE OF MAXIMUM MULTIPLICITY 2 2s 2p 1 1s

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f IRON 1s2 2s2 2p6 3s2 3p6 4s2 3d6 Orbitals are filled according to Hund’s Rule. They continue to pair up. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS HUND’S RULE OF MAXIMUM MULTIPLICITY 2 2s 2p 1 1s

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f COBALT 1s2 2s2 2p6 3s2 3p6 4s2 3d7 Orbitals are filled according to Hund’s Rule. They continue to pair up. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS HUND’S RULE OF MAXIMUM MULTIPLICITY 2 2s 2p 1 1s

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f NICKEL 1s2 2s2 2p6 3s2 3p6 4s2 3d8 Orbitals are filled according to Hund’s Rule. They continue to pair up. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS HUND’S RULE OF MAXIMUM MULTIPLICITY 2 2s 2p 1 1s

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f COPPER 1s2 2s2 2p6 3s2 3p6 4s1 3d10 One would expect the configuration of chromium atoms to end in 4s2 3d9. To achieve a more stable arrangement of lower energy, one of the 4s electrons is promoted into the 3d. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p 1 1s

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f ZINC 1s2 2s2 2p6 3s2 3p6 4s2 3d10 The electron goes into the 4s to restore its filled state and complete the 3d and 4s orbital filling. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p 1 1s

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f GALLIUM - KRYPTON The 4p orbitals are filled in exactly the same way as those in the 2p and 3p sub levels. 3 3s 3p 3d 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS HUND’S RULE OF MAXIMUM MULTIPLICITY 2 2s 2p 1 1s

THE ELECTRONIC CONFIGURATIONS OF THE FIRST 36 ELEMENTS 4 4p 4d 4f GALLIUM - KRYPTON Prefix with… 1s2 2s2 2p6 3s2 3p6 4s2 3d10 3 3s 3p 3d Ga - 4p1 Ge - 4p2 As - 4p3 Se - 4p4 Br - 4p5 Kr - 4p6 4s INCREASING ENERGY / DISTANCE FROM NUCLEUS 2 2s 2p Remember that the 4p configurations follow Hund’s Rule with the electrons remaining unpaired to give more stability. 1 1s

Practice: 1. The electronic structure of an atom of an element in Group 6 of the Periodic Table could be: A 1s2 2s2 2p2 B 1s2 2s2 2p4 C 1s2 2s2 2p6 3s2 3p6 3d6 4s2 D 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 2.

B, C, D 3. Task: work out the electronic configuration of atoms from hydrogen to argon

ELECTRONIC CONFIGURATIONS OF ELEMENTS 1-30 H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn 1s1 1s2 1s2 2s1 1s2 2s2 1s2 2s2 2p1 1s2 2s2 2p2 1s2 2s2 2p3 1s2 2s2 2p4 1s2 2s2 2p5 1s2 2s2 2p6 1s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2 1s2 2s2 2p6 3s2 3p1 1s2 2s2 2p6 3s2 3p2 1s2 2s2 2p6 3s2 3p3 1s2 2s2 2p6 3s2 3p4 1s2 2s2 2p6 3s2 3p5 1s2 2s2 2p6 3s2 3p6 1s2 2s2 2p6 3s2 3p6 4s1 1s2 2s2 2p6 3s2 3p6 4s2 1s2 2s2 2p6 3s2 3p6 4s2 3d1 1s2 2s2 2p6 3s2 3p6 4s2 3d2 1s2 2s2 2p6 3s2 3p6 4s2 3d3 1s2 2s2 2p6 3s2 3p6 4s1 3d5 1s2 2s2 2p6 3s2 3p6 4s2 3d5 1s2 2s2 2p6 3s2 3p6 4s2 3d6 1s2 2s2 2p6 3s2 3p6 4s2 3d7 1s2 2s2 2p6 3s2 3p6 4s2 3d8 1s2 2s2 2p6 3s2 3p6 4s1 3d10 1s2 2s2 2p6 3s2 3p6 4s2 3d10 ELECTRONIC CONFIGURATIONS OF ELEMENTS 1-30

ELECTRONIC CONFIGURATION OF IONS Positive ions (cations) are formed by removing electrons from atoms Negative ions (anions) are formed by adding electrons to atoms Electrons are removed first from the highest occupied orbitals (EXC. transition metals) SODIUM Na 1s2 2s2 2p6 3s1 1 electron removed from the 3s orbital Na+ 1s2 2s2 2p6 CHLORINE Cl 1s2 2s2 2p6 3s2 3p5 1 electron added to the 3p orbital Cl¯ 1s2 2s2 2p6 3s2 3p6

ELECTRONIC CONFIGURATION OF IONS FIRST ROW TRANSITION METALS Positive ions (cations) are formed by removing electrons from atoms Negative ions (anions) are formed by adding electrons to atoms Electrons are removed first from the highest occupied orbitals (EXC. transition metals) SODIUM Na 1s2 2s2 2p6 3s1 1 electron removed from the 3s orbital Na+ 1s2 2s2 2p6 CHLORINE Cl 1s2 2s2 2p6 3s2 3p5 1 electron added to the 3p orbital Cl¯ 1s2 2s2 2p6 3s2 3p6 FIRST ROW TRANSITION METALS Despite being of lower energy and being filled first, electrons in the 4s orbital are removed before any electrons in the 3d orbitals. TITANIUM Ti 1s2 2s2 2p6 3s2 3p6 4s2 3d2 Ti+ 1s2 2s2 2p6 3s2 3p6 4s1 3d2 Ti2+ 1s2 2s2 2p6 3s2 3p6 3d2 Ti3+ 1s2 2s2 2p6 3s2 3p6 3d1 Ti4+ 1s2 2s2 2p6 3s2 3p6

Electrons & orbitals Key Words: Electron density Electron cloud Electron density map Objectives: Electrons density maps and the shape & orientation of orbitals Outcomes: A-C: - Recall what electron density maps are Know & recall the shapes and orientations of orbitals: p and d

Electron configuration & chemical properties Key Words: Periodic law Groups Periods Transition metals Metalloids Lanthanides Actinides Objectives: electronic structure determines the chemical properties of an element periodic table is divided into blocks Outcomes: D: recall that chemical properties are related to electronic structure - Know the blocks of the periodic table A-C: Know the chemical properties of: s-block elements d-block elements p-block elements

Periods & groups: The reactivity of an element, and how it combines with other elements, is determined by its arrangement of electrons in its outer shell The periodic table arranges elements in order of their atomic number Groups: the vertical columns in the periodic table Periods: the horizontal rows in a periodic table

All the elements in a period have the same number of electron shells. So, the elements in each group and period show particular characteristics and trends in their chemical and physical properties Periodic Law: the properties of the elements are a function of their atomic numbers

Lanthanides + actinides Blocks Block Groups Subshell: s 1 + 2 Outer electrons in s subshells p 3+4+5+6+7+0 Outer electrons in p subshells d Transitional metals Outer electrons in d subshells f Lanthanides + actinides Outer electrons in f subshells

s-block elements Reactive metals Lower melting temperature Than other metals Reactive metals Lower melting temperature Lower boiling temperature Lower density Conduct electricity Include hydrogen and helium – but usually treated as a separate group.

d-block elements Called Transitional metals Less reactive that Group 1+ 2 metals – this is because the inner d orbital is being filled while the outer s orbital is full All conduct electricity and heat Are shiny, and hard Ductile – pulled into shape Malleable – hammered into shape Mercury is the only exception – low melting temperature liquid at room temperature

f-block elements Lanthanides – are all similar Actinides – all radioactive Only the actinides up to uranium are naturally occurring The others have all been synthesises by scientists and have extremely short half-lives

p-block elements All the non-metals and metalloids Include Tin and Lead Form positive ions Form ionic bonds with non-metals Many metals in p block do not have strong metallic characteristics All conduct heat and electricity Called post transitional metals generally unreactive

Metalloids occur in a diagonal block Mostly like non-metals Conduct electricity – but poorly Silicon and germanium are responsible for microchips Non-metals all form covalent bonds with other non-metals & ionic bonds with metals Majority do not conduct electricity Some elements form giant covalent structures

Practice:

C, B, D, B

A,

Trends in the Periodic Table Key Words: Atomic radius Ionisation energy Melting temperature Objectives: Understand trends in the periodic table Outcomes: D: understand and describe the trends in the periodic table A-C: Explain the trends in the periodic table - - - - ionization energy based on given data or recall of the shape of the plots of ionization energy versus atomic number using ideas of electronic structure and the way that electron energy levels vary across the period. - melting temperature of the elements based on given data

Key Words Ionization energy: the amount of energy it takes to strip away the first electron Electronegativity: a measure of how tightly an atom holds onto its outer shell electrons Nuclear charge: the attractive force between the positive protons in the nucleus and the negative electrons in the energy levels. The more protons, the greater the nuclear charge. Shielding: inner electrons tend to shield the outer electrons from the attractive force of the nucleus. The more energy levels between the outer electrons and the nucleus, the more shielding.

Atomic radius: measure of the size of atoms, usually measured from the nucleus to the outer shell Ionic radius: the size of ions

Key points: Across periodic table: elements gain electrons Down a group : elements gain electron shells. This changes the diameter of atoms which affects their physical and chemical properties

The atomic radius generally decreases across a period: The nuclear charge becomes increasingly positive as the number of protons in the nucleus increases. The number of electrons also increases BUT they are all in the same shell This means that they are attracted more strongly to the nucleus – so reducing the atomic radius across a period

The atomic radius generally increases down a group: The outer electrons enter new energy levels down a group So, even though the nucleus has more protons, the electrons are further away and they are screened by more electron shells. So, they are not held so tightly and the atomic radius increases

Atoms to ion: The atomic radius changes when atoms form ions Positive ions always have a smaller ionic radius that the original atom. Because: the loss of electron(s) means that the remaining electrons each have a greater share of the positive charge of the nucleus so are more tightly bound And when an ion in formed, a whole ion shell is usually lost

Negative ion has a larger atomic radius than that of the original atom even though the extra electrons are in the same electron shell, the addition of the negative charge means that the electrons are less tightly bound to the nucleus So the atomic radius is larger

Periodic Trends in Ionisation Energy: Then more tightly held the outer electrons, the higher the ionisation energy 3 main factors affecting ionisation energy of an atom: The attraction between the nucleus & the outermost electron – decreases as the distance between them increases  reducing the ionisation energy The size of the positive nuclear charge - a more positive nucleus has a greater attraction for the outer electron  so higher ionisation energy Inner shells of electrons repel the outer electron, screening or shielding it from the nucleus - the more electron shells there are between the outer electrons and the nucleus, the less firmly held the outer electron is lower ionisation energy

Ionisation energy & periods: Ionisation energy increases across a period It becomes harder to remove an electron This is because: Increasing positive nuclear charge across the period Without the addition of extra electron shells to screen the outer electrons The atomic radius gets smaller & electrons are held more firmly – so it requires more energy to make ionisation happen The end of each period is marked by the high ionisation energy of a noble gas – this is a result of a stable electronic structure & indicates their unreactive natures

This is because of subshells within each shell (b) shows that First ionisation energies do not increase smoothly across a period This is because of subshells within each shell E.g: the first ionisation energy of Be is larger than B, Mg has a larger first ionisation energy than Al – why? For Be or Mg, an electron must be removed from a full s-shell. Full subshells are particularly stable – so it requires more energy than removing a single p electron from B or Al

Nitrogen & phosphorous have unexpectedly high first ionisation energies: They both have a half-full outer p subshell. Half full subshells seem to have greater stability So requires more energy Ionisation energy decreases down a group – it becomes easier to lose an electron

Patterns in physical properties The physical properties are closely linked to the structure and bonding of atoms Melting temperature: the temperature at which the pure solid is in equilibrium with the pure liquid, at atmospheric pressure. this is affected by the packing & binding of atoms within a substance It changes as you go across a period

The relatively high melting temperatures of the metals (e. g The relatively high melting temperatures of the metals (e.g. Li, Mg, Al) are due to their metallic structure. The atoms are held tightly together is a ‘sea of electrons’ It takes a lot of energy to separate them Giant molecular structures (metalloids-silicon, carbon-in form of diamond): Strong covalent bonds between atoms which hold them tightly in a crystal structure Very difficult to remove individual atoms So very high melting temperature

Simple molecular structures: Most non-metals found on right of periodic table Small, individual molecules Strong covalent bonds within molecules But, molecules are held together by weak intermolecular forces Can be separated easily Low melting temperature

Practice D