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Flashcards by urAnonGamer Uk, updated more than 1 year ago
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| Question | Answer |
| Describe the Periodic Table in terms of arrangement of elements Explain that atoms of elements in a group have similar outer shell electron configurations resulting in similar properties | Elements in the periodic table are arranged in order of increasing atomic (proton) number Periods - horizontal rows which show repeated trends across them Group = vertical column which contain elements of similar properties. Elements in within a group have the same number of outer shell electrons, this is because they have repeating pattern of electron configuration - each element in the same group has the same number of outer electrons in the same type of orbital. This means elements in the same group will react in similar ways |
| Define periodicity | Periodicity is a regular variation of properties of elements with atomic number and position in the periodic table - the trend in properties that is repeated across each period. |
| Explain the variation of first ionisation energies and atomic radii of elements across a period | Across a period: The number of protons in the nucleus increases across a period, increasing the nuclear charge, and therefore the nuclear attraction acting on the electrons. Across a period, all electrons are added to the same shell therefore there is no increased effect of shielding to outweigh the increased nuclear attraction, therefore electron shell are drawn inwards towards the nucleus, decreasing the atomic radius. Therefore, there is a larger force of attraction felt by the outer electrons as you move across the period, resulting in a higher first ionisation energy as more energy is required to overcome it. |
| Explain the variation of first ionisation energies and atomic radii of elements across a period | However, going across the period, elements in group 3 will have a slightly smaller ionisation energy than that of group 2. This is because between these groups electrons transition from being added to the s orbital, to being added to the p orbital. The s orbital will therefore slightly increase the shielding felt by the outer p orbital electron, decreasing the ionisation energy. There is a sharp decrease in first ionisation energy between the end of one period and the start of the next as a new shell is added, which increases the distance of the outer electrons from the nucleus and increases the shielding felt by the outer electrons. These factors outweigh the increased nuclear attraction. |
| Explain the variation of first ionisation energies and atomic radii of elements down a group | Down each group the number of shells increases, so outer electrons are further from the nucleus. This also means that the effect of electron shielding increases as there are more inner shells. These factors outweigh the increased nuclear attraction from the increased number of protons. Therefore the nuclear attraction felt by the outer electrons decreases so electrons are not drawn as tightly in towards the nucleus causing the atomic radius to increase, and the first ionisation energy to decrease down a group. |
| Describe the variation in melting and boiling point for elements in periods 2 and 3. | From group 1-4 there is a general increase in boiling point across the period. Between groups 4 and 5 there is a sharp decrease. In groups 5-0 elements have low boiling points in comparison to those in groups 1-4. This variation is due to the change in structure of the elements. Groups 1-2 in period 2 and 1-3 in period 2 are giant metallic, with strong electrostatic attraction (metallic bonds) between ions and electrons The melting and boiling points increase from group 1-3 in period 3, illustrating the increasing charge on the ion and increasing number of delocalised electrons, leading to increased attractive forces. Groups 3-4 in period 2 and group 4 in period 3 are giant covalent, with strong covalent bonds between the atoms Groups 5-0 in both periods are simple molecular forces with only weak intermolecular forces between the molecules. |
| Describe the arrangement of metals and non-metals within the periodic table | ELements on the left of the periodic table are metals, with those to the very right being non-metals. However, metals which separate metals from non-metals on the periodic table, e.g silicon, are referred to as semi-metals or metalloids and have properties of both metals and non-metals. Silicon has the shiny appearance of a metal but is brittle, and can conduct electricity but only very poorly. |
| Describe the trend of electronegativity across a period | Across a period, the number of protons in the nucleus increases, therefore the nuclear attraction acting on the outer electrons will increase, as the number of shells remains the same so there is increased effect of electron shielding. The elements with greater nuclear attraction acting on their outer electrons will have a greater ability to attract electrons in a bonding pair to them, giving an increased electronegativity. Electronegativity decreases down the groups as the number of shells increases therefore increasing the effects of electon shielding and the distance of the outer electrons from the nucleus, outweighing the increased nuclear attraction from the nucleus. The most electronegative elements are Nitrogen, Oxygen and Fluorine (NOF) they are furthest across period 1 |
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