We think you have liked this presentation. If you wish to download it, please recommend it to your friends in any social system. Share buttons are a little bit lower. Thank you!
Presentation is loading. Please wait.
Published byRebecca Morrow
Modified over 3 years ago
Menzies High School AS Chemistry © KeMsoft2004
AS Module 1 Atomic Structure Fundamental particlesbe able to describe the properties of protons, neutrons and electrons in terms of relative charge and relative mass Protons, neutrons and electrons understand the importance of these particles in the structure of the atom Mass number and isotopes be able to recall the meaning of mass number (A) and atomic (proton) number (Z) be able to explain the existence of isotopes understand the principles of a simple mass spectrometer, limited to ionisation, acceleration, deflection and detection be able to interpret simple mass spectra of elements and calculate relative atomic mass from isotopic abundance, limited to mononuclear ions know that mass spectrometry can be used to determine relative molecular mass Electron arrangement be able to describe the electronic structures of atoms and ions up to Z = 36 in terms of levels s, p and d, considered as energy levels not quantum numbers understand how ionisation energies in Group II (Be - Ba) and in Period 3 (Na - Ar) give evidence for electron arrangement in levels and sub-levels © KeMsoft2004
AS Module 1 Periodicity Classification of elements in s, p, and d blocks be able to classify an element as s, p or d block according to its position in the Periodic Table Properties of the elements of Period 3 (Na - Ar) to illustrate periodic trends be able to describe the trends in atomic radius, first ionisation energy, electronegativity, melting and boiling points of the elements Na - Ar understand the reasons for the trends in these properties Group II understand the trends in atomic radius, first ionisation energy, electronegativity and melting pouint of the elements Be - Ba know the reactions of the elements Be - Ba with water and recognise the trend know the relative solubilities of the hydroxides of the elements Be - Ba and that Mg(OH)2 is sparingly soluble know the relative solubilities of the sulphates of the elements Be - Ba and that BaSO4 is insoluble and is formed in the test for sulphate ions know that beryllium is atypical, limited to covalent character (e.g. in BeCl2), the amphoteric character of Be(OH)2 and the limitation of maximum co-ordination number to four © KeMsoft2004
M+ M + e- M M+ What is IonisationThe removal of an electron(s) from a gaesous atom to form an ion. + 11+ 11- 10- 1+ M+ M (g) (g) The overall process is represented by an equation; (g) + e- M M+ © KeMsoft2004
Ionisation Energy Since the energy involved in ionising a single atom is so small we define it for... 1 mol of electrons The 1st ionisation energy of an atom is the energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of unipositive ions © KeMsoft2004
Ionisation Energy - TrendsI II III IV V VI VII O H: 1s1 1 H He He:1s2 p Li: 1s22s1 Na: 1s22s22p63s1 Be: 1s22s2 Mg: 1s22s22p63s2 B: 1s22s22px1 Al: 1s22s22p63s23p1 Li Be B C N O F Ne 2 C: 1s22s22px12py1 Si: 1s22s22p63s23px13py1 N: 1s22s22px12py12pz1 P: 1s22s22p63s23px13py13pz1 O: 1s22s22px22py12pz1 S: 1s22s22p63s23px23py13pz1 F: 1s22s22px22py22pz1 Cl: 1s22s22p63s23px23py23pz1 3 Na Mg Ne: 1 s22s22p6 Ar: 1s22s22p63s23p6 Al Si P S Cl Ar d 4 © KeMsoft2004
Ionisation Energy - TrendsI II III IV V VI VII O K: Ar4s1 Sc: Ar4s23d1 1 H He Ca: Ar4s2 Ti: Ar4s23d2 p V: Ar4s23d3 Zn: Ar4s23d10 Cr: Ar4s13d5 Ga: Ar4s23d104p1 Li Be Mn: Ar4s23d6 Ge: Ar4s23d104p2 B C N O F Ne 2 Fe: Ar4s23d7 As: Ar4s23d104p3 Co: Ar4s23d8 Se: Ar4s23d104p4 Ni: Ar4s13d10 Br: Ar4s23d104p5 3 Na Mg Cu: Ar4s13d10 Kr: Ar4s23d104p6 Al Si P S Cl Ar d K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 4 Ba Sr © KeMsoft2004
Ionisation Energy - TrendsI II III IV V VI VII O understand how ionisation energies in Group II (Be - Ba) 1 H He p and in Period 3 (Na - Ar) Li Be B C N O F Ne 2 give evidence for electron arrangement in levels and sub-levels 3 Na Mg Al Si P S Cl Ar d K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 4 Ba Sr © KeMsoft2004
Ionisation Energy - TrendsFactors affecting the size of ionisation energy The greater the attraction between the nucleus and the outer electron The greater will be the value of the 1st I.E. 1s 2s 2p 3s 3p The size of that attraction will be governed by: charge on the nucleus. The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it. distance of the electron from the nucleus. Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away. number of electrons between the outer electrons and the nucleus. If there are filled shells of electrons between the outer electron and the nucleus they will ‘cut off’ some of the attractive force of the nucleus – this known as screening or shielding. © KeMsoft2004
Ionisation Energy - Trends1. Look up and tabulate the 1st ionisation energies of the a) The Period 3 elements, b) The group II elements 2. Plot, by hand, graphs of the 1st ionisation energies of... The Period 3 elements, b) The group II elements against their atomic numbers. © KeMsoft2004
Ionisation Energy - TrendsVariation in 1st Ionisation Energy with Proton Number Proton number Ionisation energy kJ/mol P3 Ionisation Energy - Trends Variation in 1st Ionisation Energy with Proton Number Ionisation energy kJ/mol P3 Proton number © KeMsoft2004
Ionisation Energy - TrendsVariation in 1st Ionisation Energy with Proton Number Ionisation Energy - Trends Ionisation energy kJ/mol P3 Proton number 1600 Variation in ionisation energy across Period 3 1400 General trend: As the proton number increases, the 1st I.E. increases 1200 1000 1st Ionisation energy kJ/mol 800 Successive increase in nuclear charge (+ve) means the outer electrons become more strongly held and need more energy to remove them. 600 400 400 charge on the nucleus. The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it. 200 Proton number [Na Mg Al Si P S Cl Ar] © KeMsoft2004
Explain why boron has a lower first ionisation energy than beryllium.Variation in 1st Ionisation Energy with Proton Number 1s 2s 2p 3s 3p Ionisation Energy - Trends Ionisation energy kJ/mol P3 Explain why aluminium has a lower first ionisation energy than magnesium. Proton number 1600 Mg’s outer electron in an s (3s) orbital (1) Al’s outer electron is in a p (3p) orbital (1) 3p higher in energy than 3s – less energy needed to remove it(1) Variation in ionisation energy across Period 3 Two places where the general trend is not followed… 1400 1200 Explain why boron has a lower first ionisation energy than beryllium. 1000 12p 13p Be’s outer electron in an s (2s) orbital (1) B’s outer electron is in a p (2p) orbital (1) 2p higher in energy than 2s – less energy needed to remove it(1) 1st Ionisation energy kJ/mol 800 Ne3s2 600 p-orbital is higher in energy than an s orbital (further from the nucleus and screened) so the electron requires less energy to remove it; aluminium will have a lower 1st I.E. than magnesium. Even though the nuclear charge has increased this is more than off-set by the greater energy and degree of shielding of the p-orbital electron. Ne3s23px1 400 400 200 Proton number [Na /Li Mg/Be Al/B Si/C P/N S/O Cl/F Ar/Ne] © KeMsoft2004
electron lost in S is paired in a 3p orbital (1)Variation in 1st Ionisation Energy with Proton Number 1s 2s 2p 3s 3p 15p Ionisation Energy - Trends Ionisation energy kJ/mol P3 P3 Explain why sulphur has a lower first ionisation energy than phosphorus. Proton number Proton number 1600 electron lost in S is paired in a 3p orbital (1) electron lost in P is unpaired in a 3p orbital (1) repulsion between paired electrons (1) less energy needed to remove a paired electron than an unpaired one (1) Variation in ionisation energy across Period 3 Variation in ionisation energy across Period 3 Two places where the general trend is not followed… 1400 1200 Ne3s23px1py1pz1 1s 2s 2p 3s 3p 16p 1000 Ne3s23px2py1pz1 1st Ionisation energy kJ/mol 800 In both cases the electron is removed from a p-orbital (same amount of shielding in both elements) but in phosphorus the electron lost is one of the unpaired p-electrons and in sulphur it is one of the paired electrons . The paired electrons repel each other and less energy is needed to remove one of them compared with the unpaired ones; sulphur therefore has a lower 1st I.E. than phosphorus. 600 400 400 200 Proton number [Na /Li Mg/Be Al/B Si/C P/N S/O Cl/F Ar/Ne] © KeMsoft2004
As the proton number increases,1000 1s 2s 4p Be 1s 2s 2p 3s 12p 1s 2s 2p 3s 20p 3p 4s 1s 2s 2p 3s 38p 3p 4s 3d 5s Ionisation Energy - Trends Variation in ionisation energy down Group II Mg Ca Sr 1st Ionisation energy kJ/mol 800 600 400 400 Successive increase in nuclear charge (+ve) does not make up for: the increased degree of shielding felt by the outer s-electron Its higher energy state Its greater distance from the nucleus So, the outer electrons become less strongly held and need less energy to remove them. General trend: As the proton number increases, the 1st I.E. decreases 200 Proton number [Be Mg Ca Sr Ba] © KeMsoft2004
LOG10(IONISATION EERGY/kJmol) ELECTRON BEING REMOVEDCa: 1s22s22p63s23p64s2 Ca: 4s23p63s22p62s21s2 ELECTRON BEING REMOVED © KeMsoft2004
LOG10(IONISATION EERGY/kJmol) ELECTRON BEING REMOVEDCa: 4s23p63s22p62s21s2 ELECTRON BEING REMOVED © KeMsoft2004
Order of filling orbitals 1s 2s2p 3s3p3d 4s4p4d4f 5s5p5d5f 6s6p6d(6f) 7s(7p)(7d)(7f)
Atomic Structure Composition of an atom Atoms are made up of 3 fundamental subatomic particles: Relative mass Relative electric charge Position in atom.
Periodic Relationships Among the Elements
Atomic Structure Ionisation Energies. Ionisation Energy The first ionisation energy of an element is the energy required to remove completely one mole.
Trends in the Periodic Table (Chpt. 7). 1. Atomic radius (size) 2. Ionization energy 3. Electronegativity The three properties of elements whose changes.
Atomic Structure - Questions 1. What are the three sub atomic particles that make up the atom? 2. Draw a representation of the atom and labelling the sub-atomic.
Trends in the Periodic Table. Development of the Periodic Table The periodic table was invented by Dimitri Mendeleev (1869). He arranged elements in order.
IB Chemistry ATOMIC THEORY. Atomic Structure Atoms are very small ~ meters All atoms are made up of three sub-atomic particles: protons, neutrons.
The Schrödinger Model and the Periodic Table. Elementnℓms H He Li Be B C N O F Ne.
IONISATION ENERGY. WHAT IS IONISATION ENERGY? Ionisation Energy is a measure of the amount of energy needed to remove electrons from atoms. As electrons.
AS Chemistry Unit 1 Module 3 –The Periodic Table Periodicity.
The Trends in Elements in 1-20 Atomic Size Ionisation Energies.
Atomic Structure. Fundamental Particles Knowledge and understanding of atomic structure has evolved over time. Atoms are made from three types of particles:
IONISATION ENERGY OBJECTIVES: To define the term ‘ionisation energy’ To describe and explain the trends in ionisation energy across period 3 and down group.
IONISATION ENERGY A guide for A level students KNOCKHARDY PUBLISHING 2008 SPECIFICATIONS.
Learning Objectives: Define first ionisation energy and successive ionisation energy. Explain the factors that influence ionisation energies. Predict the.
Homework Private study work (bring notes to show me next lesson); Read pages 40 – 41 in your text book and complete the practice questions on each double.
Periodic Table Trends. Metallic character Metallic Character Depends on how readily the element gives up their valence e-’s.
Atom atom atom atom atom 1.True or false? Protons are in the nucleus.
Title: Lesson 7 Successive and First Ionisation Energies Learning Objectives: Understand why different elements have different ionisation energies Know.
General Chemistry Principles & Modern Applications 9 th Edition Petrucci/Harwood/Herring/Madura Chapter 9 The Periodic Table and Some Atomic Properties.
Periodicity Physical Properties Ionisation energies Li Rb Kr K Ar Na Ne He.
CHAPTER 8 ELECTRON CONFIGURATIONS AND PERIODICITY.
Relative energy levels of electrons in gaseous atoms of the first twenty elements Increasing energy s p d f 1s Electronic Structure Energy levels within.
Electronic Configuration Ochran THE BOHR MODEL OF THE ATOM A small nucleus of protons and neutrons surrounded by electrons in shells each shell.
ATOMIC NUMBER 1st IONISATION ENERGY / kJmol -1 Variation in 1st Ionisation Energy EXPLANATION Despite having a nuclear charge of only 1+, Hydrogen has.
Ch 5.3 Electron Configuration and Periodic Properties.
1 Periodic Table II Periodic table arranged according to electron arrangement Periodic table also arranged according to properties? Properties must depend.
IONISATION ENERGY CONTENTS What is Ionisation Energy? Definition of 1st Ionisation Energy What affects Ionisation Energy? General variation across periods.
Atomic Structure. Simple model of an atom An atom is made of a tiny nucleus with electrons orbiting around it. The nucleus is made up of protons and neutrons.
Atomic structure Greeks Rutherford GCSE ATOM: Electrons in shells Ionisation Energy: Experimental evidence A level Atom and the existence of orbitals.
Lesson objectives Define first ionisation energy and successive ionisation energy. Explain the factors that influence ionisation energies. Predict the.
5: Trends in the periodic table j.represent data, in a graphical form, for elements 1 to 36 and use this to explain the meaning of the term ‘periodic property’
Revising Atoms. Learning Objectives Candidates should be able to: Identify and describe protons, neutrons and electrons in terms of their relative charges.
Trends in the periodic table. Atomic radius Atomic radii trends and explanations Atomic radius decreases across a period because each successive element.
Chapter 3Atoms and Elements 3.6 Isotopes and Atomic Mass 1 24 Mg 25 Mg 26 Mg Copyright © 2009 by Pearson Education, Inc.
Lecture 1.1 Refresh your high school chemistry CS882, Fall 2006.
New Way Chemistry for Hong Kong A-Level Book 1 1 Chapter 5 Electronic Configurations and the Periodic Table 5.1 Relative Energies of Orbitals 5.2 Electronic.
© AS Jul-12. Electronegativity = the power of an atom to attract the electrons in a covalent bond.
Drill – 11/19 What is meant by “periodic trend”?.
Quiz Grades Rewrite your answer to question 1 on the back of your quiz. If it is correct, I will add up to 20% to your quiz grade (maximum score: 100%).
AS Chemistry The Periodic Table. THE PERIODIC TABLE occur in the same vertical columns (GROUPS). = the elements arranged in rows (PERIODS) such that chemically.
Unit F321 Module Electron Structure. Atomic Structure Protons, neutrons, electrons How to make ions Relative atomic mass.
PERIODIC TRENDS. CONTENTS Introduction Electron configuration Bonding & structure Atomic radius 1st Ionisation Energy Electronegativity PERIODICITY.
Atomic Structure, Bonding and Periodicity. Contents Atomic Structure Amount of Substance Bonding Periodicity.
Chemistry SASA Qtr 1 Review Guide. 1. Deuterium ( 2 H) and protium ( 1 H) are two isotopes of hydrogen. Which of the following statements best compares.
© 2017 SlidePlayer.com Inc. All rights reserved.