Trends in the Periodic Table PDF

Trends in the Periodic Table Trends in Atomic Radii **Definition: The atomic radius (covalent 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.**...

Trends in the Periodic Table Trends in Atomic Radii **Definition: The atomic radius (covalent 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.** (1) The values of the atomic radius increase down the groups in the Periodic Table for two reasons. - The additional electrons are going into a new energy level which is further from the nucleus. Since the outer electrons are becoming further away from the nucleus, the atomic radius increases. - With these extra energy levels further from the nucleus as we move down the periodic table, the screening effect of the inner electrons reduces the pull the positive nuclear charge has on the outer electrons. The values of the atomic radius decrease across a period in the Periodic Table for two reasons. - Increase in effective nuclear charge. The increased nuclear charge tends to draw the energy levels closer to the nucleus. As the attractive force increases from left to right across the Periodic Table, the nucleus pulls the outer electrons closer to it, the atoms become smaller and the atomic radius decreases. - No increase in the screening effect. All elements in the same period have the same outer energy level Trends in Ionisation Energy **Definition: The first ionisation energy of an atom is the minimum energy required to completely remove the most loosely bound electron from a neutral gaseous atom in its ground state.** 1 The values of the first ionisation energy decrease down a group in the Periodic Table for two reasons: - Increasing atomic radius. The outermost electrons are becoming further away from the attractive force of the nucleus. It becomes easier to remove an electron from the outer energy level, so the ionisation values decrease. - Screening effect of inner electrons. The outermost electrons are somewhat shielded from the attractive force of the positively charged nucleus and are easier to remove so the ionisation energy values decrease. The values of the first ionisation energy across a period in the Periodic Table for two reasons: - Increasing effective nuclear charge. The most loosely bound electron gets pulled more strongly by the nucleus as we go across the group. More energy is then needed to remove this electron, so the ionisation energy values increase. - Decreasing atomic radius. The atomic radius decreases from left to right therefore the electron in the outermost level is becoming closer to the nucleus. Increased attraction between the nucleus and the electron makes it more difficult to remove the electron so the ionisation energy value increases. Exceptions to the General Trend across a Period (2) In the second period of the table (Li to Ne) the ionisation energy increases, however there is a dip between Beryllium and Boron. Be: 1s2 2s2 B: 1s2 2s2 2p1 2 Atoms whose outermost sublevel is half full or completely full have extra stability. The outermost electron of Beryllium is completely full while the outermost shell of Boron is only half full. This means the ionisation energy of Boron is less than that of Beryllium because it takes less energy to remove an electron. The same can be seen between Nitrogen and Oxygen. N: 1s2 2s2 2p3 O: 1s2 2s2 2p4 Nitrogen is more stable because it has a half-filled outer-shell while Oxygen does not have a half-filled outer-shell. This means that more energy is required to remove an electron from Nitrogen, so the ionisation energy is higher. Further Evidence for the Existence of Energy Levels **Definition: Second Ionisation Energy is the energy required to remove an electron from an ion with one positive charge in the gaseous state i.e., the energy required to carry out the following X-(g) → X2+(g) + e-. ** This means the energy needed to remove the second loosest electron, after the loosest has already been removed. This definition can keep going to define the third, fourth, etc., ionisation energies. (3) K: 1s2, 2s2, 2p6, 3s2, 3p6,4s1 3 - n=4: One electron removed from potassium. First ionisation energy is the lowest because the electron in the 4s sublevel is the easiest to remove. - n=3: Filled outer sublevel requires more energy to remove an electron from it. The increase in the second ionisation energy suggests that these electrons are closer to the nucleus and therefore experience less shielding. This is good evidence for the existence of a lower energy level. - n=2: Since an electron is being removed from the filled 2p6 sublevel, a lot of energy is needed as this full sublevel has extra stability. - n=1: Large increase because we are removing an electron from the filled n=1 energy level which is closer to the nucleus. There are 2 electrons in the n=1 energy level. Trends in Electronegativity The values of electronegativity decreases down the groups in the Periodic Table for two reasons: - Increasing atomic radius. This means that as you go down a group, the atom gets bigger, meaning that the atom has a weaker pull on electrons in a bond as they are further away. - The Screening Effect of the inner electrons increases going down a group. This reduces the pull that the nucleus has on the outer electrons involved in bonding. The values of electronegativity increases across the periods in the Periodic Table for two reasons: - Increasing effective nuclear charge. The attraction between the nucleus and the outer electrons is steadily increasing. The electrons involved in bonding are being more strongly attracted to the nucleus so the electronegativity increases. - Decreasing atomic radius. The electrons in the outermost level are becoming closer to the nucleus. There is a greater attraction between the nucleus and these outer electrons so the electronegativity increases. Trends within Groups The chemical properties of elements are highly determined by the number of electrons in the outermost energy level. I - Alkali metals. II - Alkaline earth metals. III - Rare earth metals. IV - Crystallogens. V - Pnictogens. VI- Chalcogens. VII - Halogens. VII - Noble gases 4 Trends in Chemical Reactivity of Alkali Metals Properties - Very reactive elements because they all have low first ionisation energy values. - Do not occur free in nature, have to be extracted from their compounds. - Readily form ionic compounds by losing their single outer electron. - The reactivity of the alkali metals increases down the group because the increasing effect of inner shells of electrons cause the first ionisation energy to decrease down the group and so the electron in the outermost energy level is more easily lost. Important Reactions with Alkali Metals Reaction with Oxygen (O2) All Alkali Metals react with oxygen to form oxides: 2K + ½O2 → K2O Potassium + Oxygen → Potassium Oxide Any other Alkali Metal can replace the Potassium in this reaction. Reaction with Water (H2O) All Alkali Metals react with water to form the hydroxide of the metal. Hydrogen gas is given off: Na + H2O → NaOH + ½ H2 Sodium + Water → Sodium Hydroxide + Hydrogen Any other Alkali Metal can replace the Sodium in this reaction. - Alkali Metals are stored under oil so they won’t react with oxygen in the air - Alkali Metals lose their shine when exposed to Air - Moving down the group, the Alkali Metal become more reactive with Water Trends in Chemical Reactivity of Halogens Properties - Most electronegative elements in the periodic table - Electronegativity values decrease down the table, so fluorine is the most electronegative element of the halogens - Quite reactive as they have a great attraction for electrons - Do not exist free in nature and must be extracted from their compounds - The trend in reactivity is in the opposite direction to that of the Alkali Metals. - Remove electrons easily from other substances due to their high electronegativities 5 Exam Questions (b) Define first ionisation energy. The minimum energy required to remove the most loosely-bound electron from an isolated gaseous atom in its ground state. Explain why the first ionisation energy value of silicon is (i) greater than that of aluminium, Greater nuclear charge smaller atomic radius (ii) less than that of carbon. Greater atomic radius. (c) The successive ionisation energies of silicon are shown in the graph on the right. Explain how the graph provides evidence for energy levels in the silicon atom. Sharp increase in ionisation energy for removal of 5th electron as it is the first to be removed from the 2nd shell. Sharp increase in ionisation energy for removal of 13th electron as it is the first to be removed from the 1st shell. Gradual increase in ionisation energies for the first 4 electrons (same shell), 5th to 12th electrons (same shell) and 13th to 14th electrons (same shell) What other experimental evidence do we have for the existence of energy levels in atoms? Line Emission Spectrum 2013 – HL- Section B – Question 11 (a) Define first ionisation energy. Minimum energy required to remove the most loosely bound electron from a gaseous atom in its ground state. The graph shows the first ionisation energy values, displayed in order of increasing atomic number, for the first 31 elements of the periodic table. Refer to the table of first ionisation energy values on page 80 of the formulae and tables booklet 6 (i) Name the elements labelled B and P in the graph. B is helium and P is Sulphur What is the numerical value of x? 900 (ii) What is the principal reason for the large decrease in first ionisation energy between the elements labelled R and S? R has a full outer lower sub-level (iii) Explain why the first ionisation energy value of the element labelled H is lower than that of the element labelled G H has lower first ionisation energy because it has a less stable electron configuration than G. Half filled 2p sublevel G has higher first ionisation energy because it has a more stable electron configuration than H. Half-full 2p-sublevel is extra stable. 2012 – HL – Section B – Question 10 (a) Define electronegativity. Relative attraction that an atom of an element has for a shared pair of electrons of electrons in a covalent bond State two factors that cause electronegativity values to increase across a period in the periodic table of the elements. Increasing effective nuclear charge. Decreasing atomic radius State which of the following compounds contain intermolecular hydrogen bonds: (i) hydrogen chloride, HCl, (ii) water, H2O, (iii) ammonia, NH3. Justify your answer. H2O and NH3. In both of these molecules, hydrogen is bonded to a small highly electronegative element. Suggest a reason why the boiling point of ammonia (–33 ºC) is significantly lower than that of water (100 ºC) There is weaker and less effective hydrogen bonding in ammonia 7  References 1. Socratic.org 2. Themadscientist.net 3. Brainly.com 8