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AS inorganic Defs
All inorganic definitions for AQA AS chemistry year 12
| Term | Definition |
|---|---|
| Periodicity | Repeating patterns or trends in the chemical and physical properties of elements across a period. Often due to different electron arrangement, energy levels and sub-levels |
| Atomic radius | The distance from the nucleus to the outermost electron. This decreases as the nuclear charge increases due to greater attraction on the outer shell and increases with new electron shells due to greater electron shielding which counteracts the attraction |
| Ionisation energy | The energy required to remove one mole of electrons from one mole of gaseous substance to create a mole of gaseous substance with a positive charge. Decreases as atomic radius increases |
| Electronegativity | The ability of an atom to attract a bonding pair of electrons between itself and the atom it’s bound to |
| Reactivity | The tendency of a substance to undergo a chemical reaction |
| Periodic trends: atomic radius | Generally decreases along the period but increases down the group |
| Periodic trends: ionisation energy | Generally increases along the period but decreases down the group |
| Periodic trends: electronegativity | Generally increases along the period but decreases down the group |
| Periodic trends: reactivity | Generally metals become more reactive down the group and non-metals become more reactive along the period |
| Energy levels | Regions or shells around the nucleus where electrons are likely to be found. They’re represented by whole numbers with the lowest energy closest to the nucleus and the highest the furthest away |
| Sub levels | Within an energy level there are sub levels to represent the different shapes of electron orbitals and how many electron pairs fit on it. They are labelled s, p, d, and f |
| s block | Elements where their outermost electron is in the s sub level. These are elements with only 1 or 2 valence electrons on the left side of the periodic table |
| p block | Elements where their outermost electron is in the p sub level. They can hold up to 6 electrons |
| d block | Elements where their outermost electron is in the d sub level. They can hold up to 10 electrons |
| f block | Elements where their outermost electron is in the f sub level. They can hold up to 14 electrons |
| Period 3 trends: sodium to argon | Atomic radius decreases while ionisation energy increases. Melting points increase from sodium to silicon before sloping back down. The reactivity decreases along the period |
| Group 2 (alkali) metal trends: beryllium to radium | Atomic radius increases down the group while ionisation energy decreases. Reactivity increases down the group due to a wider atomic radius so it’s easier to lose electrons. Melting points decrease since more electron shielding means weaker metallic bonds |
| Alkali metal uses | Extracts halogens from metal halides, reducing agent, antacids, neutralising soil for agriculture, removes sulphur from flue glass, barium sulphate is insoluble and opaque to x-rays so can identify lower GI issues, barium chloride tests for sulphate ions |
| Trends of water solubility in group 2 compounds | Hydroxides and chlorides increases as it goes down the group whereas sulphate decrease in solubility |
| Halogens | Group 7 non-metals that have 7 valence electrons & are extremely reactive. Are powerful oxidising agents but their ions are powerful reducing agents |
| Halogen trends in electronegativity | Decreases down group due to increased electron shielding. As it gets larger it’s outer shell is further from the positive nucleus so there is less attraction acting on the binding pair of electrons |
| Halogen trend in boiling points | Increases down the group due to a higher electron density leading to stronger Van Der Waals forces which takes more energy to overcome |
| Halogen displacement trends | A halogen higher up in the group (the stronger oxidiser) will displace the halogen lower in the group to form a more stable molecule |
| Halide trends in reducing ability | Halide ions increase reducibility down the group due to greater electron shielding. Since the outer shell is further from the nucleus there is less attraction on those electrons making them easier to be donated |
| Silver nitrate test | Forms coloured precipitate for different halide ions: iodide forms a yellow precipitate soluble in dilute ammonia, bromide forms a cream precipitate soluble in conc. ammonia, chloride forms an insoluble white precipitate |
| Why silver nitrate is acidified | To prevent formation of metal hydroxides & carbonates, reduces solubility of silver halides in water, & minimises interference from impurities |
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