Chpt. 4: The Periodic Table

1. Chpt. 4: The Periodic Table

2. In this Chapter............ • What is an element??? • Chemists and order among the elements • The Periodic Table • Periodic Table & arrangement of electrons in • atoms

3. Robert Boyle from Co. Waterford • first chemist to define the meaning of the word • element An element is a chemical substance that cannot be split up into simpler substances by chemical means.

4. History of the Elements 1. The Greeks (400 B.C.) • The Greeks believed that • there were four • elements from which • everything else was • made: • - earth • - air • - water • - fire

5. 2. Robert Boyle (1661) • Boyle described elements as • “primitive and simple • substances”. • This is close to the modern • definition: • “An element is a substance • that cannot be split up into • simpler substances by • chemical means • Predicted that compounds • were made from and could be • broken down into elements

6. 3. Humphry Davy (early 1800’s): • Used electricity to • separate compounds • into their elements • He discovered • potassium, sodium, • calcium, barium, • strontium and • magnesium

8. 4. Henry Moseley (1914): • Using X-Rays he • discovered a method of • determining the • number of protons in • the nucleus of an atom. • The atoms of each • element have different • numbers of protons in • them i.e. no two • elements have the • same number of • protons No. of protons = Atomic No.

9. History of the Periodic Table After 1800’s more and more elements were being discovered and it was becoming increasingly difficult to understand and memorise all the properties of each element. Four men were responsible for bringing order among the elements in the form of the Periodic Table: - Johann Dobereiner - John Newlands - Dmitri Mendeleev - Henry Moseley

10. The Periodic table was designed to classify • elements according to their properties and to show • trends in their physical and chemical properties • The main groups are: I – Alkali metals; II – Alkaline • earth metals; VII – Halogens; VIII (also called group • 0) – Noble/inert gases. • The elements between groups II and III are called • the D-block or transition metals.

11. 1. Dobereiner’s Triads (1829): • Studied the properties of • various elements in a bid to • find order • Discovered that when looking • at some groups of three • elements they had similar • chemical properties and the • atomic weight of middle • element was average of the • other two elements • Group of three similar • elements called a triad

12. Dobereiner’s Triads A triad is a group of three elements with similar chemical properties in which the atomic weight of the middle element is approximately equal to the average of the other two.

13. 2. Newlands Octaves (1864): • John Newland • arranged known • elements in • order of increasing • atomic weights. • Noticed properties • seemed to repeat every • eight elements.

14. Newlands Law of Octaves An octave is a group of elements arranged in order of increasing atomic weight, in which the first and the eighth element of each group have similar properties

15. Problems with Newlands Octaves • The properties repeat every 8 as noble gases hadn’t • been discovered yet! • There were several problems such as iron being • grouped with oxygen and sulphur. • Newlands tried to force all of the known elements • into the table instead of leaving gaps for elements not • yet discovered example: Li and Ag in same group.

16. 3. Mendeleev’s Periodic Table (1869): Dmitri Mendeleev listed all known elements (63) in order of increasing atomic weight and again noticed similar properties on every eight element. Mendeleev left gaps and predicted elements that had not been discovered.

17. Mendeleev’s Periodic Law When elements are arranged in order of increasing atomic weight (relative atomic mass), the properties of the elements vary periodically.

18. Mendeleev: • What he did: • Put elements with the same properties in the same • vertical group. • Left gaps to make the elements fit into the proper • column (group). • Predicted that elements (eg. Germanium and • Gallium) would be discovered to fill these gaps. • Predicted their properties. • Reversed the order of some elements so that their • properties matched their group e.g. Te and I

19. Mendeleev's Periodic Table

20. Mendeleev: • Gained acceptance when: • new elements discovered fitted properties predicted • reversals justified by discovery of atomic number

21. 4. Moseley (1913): • Moseley studied the frequencies of the • x-rays emitted by atoms of different • elements. • Found frequencies varied depending on • the amount of positive change. • , • In other words the difference between • the elements is the number of protons • in the nucleus • Discovered method for determining • number of protons in nucleus of an • atom – known as atomic number

22. Atomic Number The atomic number of a atom is the number of protons in the nucleus of that atom *Note: Once the atomic number was known it was seen that Mendeleev’s table was in order of increasing atomic number.

23. The Modern Periodic Table The modern Periodic Table is an arrangement of elements in order of increasing atomic number The Periodic Law When elements are arranged in order of increasing atomic number, the properties vary periodically

25. Attention: New Additions to Periodic Table WOMANIUM (WO) Physical properties: Generally soft and round in form. Boils at nothing and may freeze any time. Very bitter if not used well. Chemical properties: Very active and highly unstable. Possesses strong affinity with gold, silver, platinum, and precious stones. Violent when left alone. Turns slightly green when placed next to a better specimen. Usage: An extremely good catalyst for dispersion of wealth. Caution: Highly explosive in inexperienced hands! MANIUM (XY) Physical properties: Solid at room temperature but gets bent out of shape easily. Difficult to find a pure sample. Due to rust, aging samples are unable to conduct electricity as easily as young samples. Chemical properties: Attempts to bond with WO any chance it can get. Also tends to form strong bonds with itself. Becomes explosive when mixed with Childrium for prolonged period of time. Usage: Possibly good methane source. Caution: In the absence of WO, this element rapidly decomposes and begins to smell.

26. Uses of the Periodic Table • Obtaining atomic numbers and mass numbers • Obtaining relative atomic masses • Writing electronic configurations.

27. Atomic Number & Atomic Mass Number • Atomic Number (Z): • number of protons an element has • smaller of the two numbers • Atomic Mass Number (A): • - of an element is the sum of the number of • protons and neutrons in the nucleus of an • atom of that element • - larger of the two numbers • - unit = atomic mass unit a.m.u. • 1 a.m.u. = 1.66 x 10-24g • - always a whole number

28. Nuclear Formula of an element

29. Using the Nuclear Formula – Sample Questions • How many protons, neutrons and electrons are in the following: • i) 3517Cl • ii) 4521Sc3+

30. Mass Spectrometry 1919 an instrument called mass spectrometer was built to measure the mass of atoms.

31. Mass Spectrometer

32. Isotopes Isotopes are atoms of the same element (i.e. they have the same atomic number) that have different mass numbers due to different numbers of neutrons in the nucleus. Since the neutron has no charge, the quantity of neutrons in an atom can change slightly without having an effect on the atom. Example: Neon – 2 isotopes, Chlorine – 2 isotopes, Carbon – 3 isotopes and Hydrogen – 3 isotopes

33. 3 Isotopes of Hydrogen

34. 3 Isotopes of Carbon

35. Calculating the average mass of an atom: A sample of chlorine is found to consist of 75% 3517Cl and 25% 3717Cl . Calculate the average mass of an atom of chlorine. Average Mass = (Mass isotope 1 x % abu) + (Mass isotope 2 x % abu) 100

36. Relative Atomic Mass (Atomic Weight) Ar The relative atomic mass (Ar) of an element: - is the average of the mass numbers of the isotopes of the element, - as they occur naturally, - taking their abundances into account and, - expressed on a scale in which the atoms of the carbon-12 isotope have a mass of exactly 12 units. (average mass of an atom, measured relative to the mass of the carbon -12 isotope)

37. The carbon-12 isotope has a perfect mass of 12amu. • 1/12 of this mass is a perfect 1.000 • All average masses represented on the periodic table • are compared to this value of 1.000. We therefore say • that these atomic masses are relative to the mass of • 1/12th of the C-12 isotope. • Ratio: • Ar= mass of atom of element 1/12 mass of atom of carbon-12 Note: ratio therefore no units

38. Relative Atomic Mass Calculations Formula: Ar= (Mass isotope 1 x % abu) + (Mass isotope 2 x % abu) 100

39. Example 1: An element, X, consists of 92.2% atoms with a mass 28, 4.7% of atoms with a mass 29 and 3.1% of atoms with a mass 30. What is the relative atomic mass? What is the element?

40. Example 2: The two isotopes of chlorine have mass numbers of 35 and 37 respectively. Taking the relative atomic mass of chlorine to be 35.46, calculate the % of each isotope present in the element

41. Writing Electronic Configurations • is the arrangement of electrons in an atom • two methods: • - Bohr Model – in terms of energy levels (simple) • - Energy sublevels

42. Bohr Model Method: • Find atomic number of element. • As all elements are neutral - number of protons equals • number of electrons. • Remember: • n = 1 energy level holds 2 electrons • n = 2 energy level holds 8 electrons • n = 3 energy level holds 8 electrons • n = 4 energy level holds 18 electrons

43. Bohr Model Example: Write the electronic configuration for potassium showing the number of electrons in each main energy level. From periodic table potassium has atomic number 19 therefore since neutral atom: no. of protons = no. of electrons = 19 1st energy level = 2 e- 2nd energy level = 8 e- 3rd energy level = 8 e- 4th energy level = 1 e- So, electronic configuration of potassium is (2,8,8,1)

44. Try the following: Write the electronic configuration for the following showing the number of electrons in each main energy level: i) Fluorine ii) Calcium Note: - you must be able to write the electronic configurations of the first 20 elements in terms of the number of electrons in each main energy level. - number of electrons in outer shell is same as group number e.g. Lithium group I has one outer electron, Boron group III has three outer electrons

47. Electron Configuration - in terms of sublevels Remember: Main energy level Sublevel

48. Electronic Configuration – in terms of sublevels An energy sublevel is a group of orbitals, within an atom, all having the same energy. One orbital in an ‘s’ energy sublevel Three orbitalsin a ‘p’ energy sublevel - X, Y, Z Five orbitalsin a ‘d’ energy sublevel In terms of energy s<p<d<f. Electrons will fill sublevels in order of increasing energy.

49. S sublevel (orbital) can hold 2 electrons P sublevel can hold 6 electrons, 2 in each px, py and pz orbital d sublevel can hold 10 electrons, 2 in each of the 5d orbitals

50. Main Energy Level – 1, 2, 3, 4 etc. Sublevel – s, p, d, f Orbitals – 1s, 2s, 2px etc. Electrons – 2 in each orbital