Definitions

1. Definitions

2. Energy level • is the fixed energy value that an electron in an atom may have. Exam Q (Hons) ‘08/Q10c ‘07/Q4

3. Ground State • lowest energy state ( in 1s orbital) • Excited state = higher energy state Exam Q (Hons) 08/Q10c

4. An orbital • is a region in space within which there is a high probability of finding an electron. Exam Q (Hons) ‘06/Q5

5. An element • is a substance that cannot be split up into simpler substances by chemical means.

6. 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. (Dobereiner)

7. Newlands’ Octaves • are groups of elements arranged in order of increasing atomic weight, in which the first and the eighth element of each group have similar properties.

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

9. The atomic number(Z) • is the number of protons in the nucleus of that atom.

10. Periodic Table • is an arrangement of elements in order of increasing atomic number.

11. Elements are arranged • in order of increasing atomic number, the properties of the elements vary periodically.

12. Mass number (A) • is the sum of the number of protons and neutrons in the nucleus of an atom of that element.

13. Isotopes Exam Q (Hons) ’06/Q10a • are atoms of the same element ( i.e. they have the same atomic number) that have different mass numbers due to the different number of neutrons in the nucleus.

14. Relative Atomic Mass • is the average of the mass numbers of the isotopes of the element • as they occur naturally • taking their abundances into account • relative to 1/12th mass of carbon 12 atom (expressed on a scale in which the atoms of carbon 12 isotope have a mass of exactly 12 units). Exam Q (Hons) ’06/Q10a

15. how? why? A+ Victor Mass Spec. • V • I • A • S • D Vaporisation Ionisation Acceleration Separation Detection

16. Aufbau Principle • that when building up the electronic configuration of an atom in its ground state, the electrons occupy the lowest available energy level.

17. Hund’s Rule of Maximum Multiplicity • states that when two or more orbitals of equal energy are available, the electrons occupy them singly first before filling them in pairs.

18. Pauli Exclusion Principle • that no more than two electrons may occupy an orbital and they must have opposite spins.

19. Compound • is a substance that is made up of two or more different elements combined together chemically.

20. Octet Rule • that when bonding occurs, atoms tend to reach an electron arrangement with eight electrons in the outermost shell.

21. An Ion • is a charged atom or group of atoms.

22. An Ionic bond • is the force of attraction between oppositely charged ions in a compound.

23. A transition metal • is one that forms at least one ion with a partially filled d sublevel.

24. Molecule • is a group of atoms joined together. It is the smallest particle of an element or compound that can exist independently.

25. Valency of an element • is defined as the number of atoms of hydrogen or any other monovalent element with which each atom of the element combines.

26. Electronegativity • is a measure of the relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond. Exam Q (Hons) ’06/Q5

27. Electronegativity • difference > 1.7 indicates ionic bonding in a compound. • An electronegativity difference ≤ 1.7 indicates covalent bonding in a compound.

28. The value of electronegativity • decrease down the groups in the Periodic Table for two reasons: • increasing atomic radius • screening effect of inner electrons

29. The values of electronegativity • increase across the periods in the Periodic Table for two reasons: • increasing nuclear charge • decreasing atomic radius F= most electronegative element. Halogens –decrease in reducing power down the group due to drop in electroneg. values.

30. Intermolecular Forces • attractive (repulsive) forces between molecules • Intramolecular forces are attractive (repulsive) forces within a molecule

31. Vans der Waals Forces • are weak attractive forces between molecules resulting from the formation of temporary dipoles.

32. Dipole-dipole • Dipole – dipole forces are forces of attraction between the negative pole of one molecule and the positive pole of another.

33. Hydrogen bonds • are particular types of dipole-dipole attractions between molecules in which hydrogen atoms are bonded to nitrogen, oxygen or fluorine. • The hydrogen atom carries a partial positive charge and is attracted to the electronegative atom in another molecule. Thus, H acts as a bridge between two electronegative atoms.

34. The Law of Conservation of Mass • the total mass of the products of a chemical reaction is the same as the total mass of the reactants.

35. The Law of Conservation of Matter • that in any chemical reaction, matter is neither created nor destroyed but merely changes from one form into another.

36. Tests for Anions • Chloride • Sulfate/sulfite • carbonate/hydrogen carbonate • nitrate • phosphate • (NB know confirmatory test too!)

37. Chloride • Add AgNO3 • Get white ppt • Confirm = ppt dissolves in dilute ammonia • Equation needed

38. Sulfate/sulfite • Add BaCl2 • Get white ppt • Distinguish • add dil HCl to white ppt • ppt remains = sulfate • ppt dissolves = sulfite Equation needed !!

39. CO32- /HCO3- • Add dil. HCl (or any acid) • Get bubbles of CO2 (limewater milky) • Distinguish • add MgSO4 to fresh solution • Get white ppt. immediately = carbonate white ppt on heating = hydrogen carbonate Equation needed !!

40. Nitrate • Brown Ring Test • Add fresh FeSO4 At slant add conc. H2SO4 drop wise • Get brown ring at junction of 2 layers No equation needed

41. Phosphate • Add ammonium molybdate • Add 5 drops of conc. nitric acid (warm the solution) • Get yellow ppt No equation needed Confirm: Goes colourless when add dilute NH3

42. The atomic radius of an atom • is defined as half the distance between the nuclei of two atoms of the same element that are joined together by a single covalent bond. Exam Q (Hons) 07/Q4

43. The values of atomic radius • increase down any one group in the Periodic Table for two reasons: • extra shell • screening effect of inner electrons

44. The values of atomic radius • decrease from left to right across a Periodic Table for two reasons: • increasing nuclear charge • no increase in screening effect

45. The first ionisation energy of an atom • is the minimum energy required to completely remove the most loosely bound electron from one mole of neutral gaseous atom in the ground state.****** • 2004 =9 marks (2.25%) • 2002 = 8marks(2%)

46. The values of ionisation energy • decrease down the groups in the Periodic Table for two reasons: • increasing atomic radius • screening effect of inner electrons

47. The values of ionisation energy • increase across the Periodic Table for two reasons: • increasing nuclear charge • decreasing atomic radius

48. More on ionisation energy • First Ionisation Energy • M – e-M+ • Second Ionisation energy M+ – e-M2+ Major jump in I.E. values – significance

49. The value of electronegativity • decrease down the groups in the Periodic Table for two reasons: • increasing atomic radius • screening effect of inner electrons

50. The values of electronegativity • increase across the periods in the Periodic Table for two reasons: • increasing nuclear charge • decreasing atomic radius F= most electronegative element. Halogens –decrease in reducing power down the group due to drop in electroneg. values.