7291576
Description
Flashcards by Ella Broomer, updated more than 1 year ago
|
Created by Ella Broomer
almost 8 years ago
|
|
| Question | Answer |
| Which is the mass number and which is the atomic number? | |
| Is the top or along the edge the period or the group? | Period (down the peridoic table) - the number of the highest energy electron shell in element's atoms. Group (across the peridoic table)- same number of outer-shell electrons and similar properties |
| Periodicity (repeating trends in properties of elements) need to know about | Electron configuration Ionisation energy Structure Melting point |
| Periodic trend in electron configuration | |
| Advantages of groups 1-7 and 0 (old way) and not 1-18 | (Old way) Group number matches number of electrons in highest energy electron shell. 1-18 does not match. The periodic table in the exam has both |
| First ionisation energy | (How easily an atom loses electrons to form positive ions) Energy required to remove one electron from from each atom in one mole of gaseous ions of an element to form one mole of gaseous +1 ions |
| Ionisation energy | Greater distance between nucleus and outer electrons, less nuclear attraction 1st electron lost is in highest energy level, experience least attraction from nucleus |
| Factors affecting attraction and therefore ionisation energy- Atomic Radius | Greater distance from nucleus and outer electrons = less nuclear attraction Force of attraction falls off sharply with distance |
| Factors affecting attraction and therefore ionisation energy- Nuclear charge | More protons in nucleus = greater attraction between nucleus and outer electrons |
| Factors affecting attraction and therefore ionisation energy- Electron shielding | Inner-shell electrons repel outer-shell electrons (as electrons are negatively charged) The repulsion called the shielding effect, reduces attraction between nucleus and outer electrons |
| Second ionisation energy | Energy required to remove one electron from each atom in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions |
| Describe the difference between the size of first and second isonisation energy e.g. of helium | Second ionisation energy is greater than first Helium has 2 protons attracted to 2 electrons in 1s sub shell After 1st electron is lost the 2nd is pulled closer to the nucelus Nuclear attraction on remaining electrons increases, more ionisation energy needed to remove 2nd electron |
| Describe the successive ionisation energy graph of sodium | Large increase between electron 9 and 10 suggesting 10th electron is removed from a different shell. Large increase between 1 and 2 suggesting 2nd electron is removed from different shell Remember n=1 is electrons 9 and 10 |
| Successive ionisation energies allow predictions to be made about: | Number of electrons in the outer shell Group of the element in the periodic table Identity of the element |
| From the first 4 ionisation energies identify the element from period 3 | Steady increase Large increase between 3rd and 4th ionisation energies 4th electron must be removed from inner shell 3 electrons in the outer shell so must be in group 3 Period 3 and Group 3 (13) so must be aluminium |
| Key patterns with first ionisation energies | General increase in 1st ionisation energy across each period e.g. H → He, Li → Ne, Na → Ar Sharp decrease in 1st ionisation energy between the end of 1 period and start of the next e.g. He → Li, Ne → Na, Ar → K Trends are due to atomic radius, electron shielding and nuclear charge |
| Why does 1st ionisation energy show an increase across a period? | Nuclear charge increase, same shell: similar shielding, nuclear attraction increase, atomic radius decrease, 1st ionisation increases |
| Why does 1st ionisation energy show a decrease down a group? | Atomic radius increases, more inner shells so shielding increases, nuclear attraction on outer electrons decreases, 1st ionisation energy decreases |
| Why does nuclear charge increase but 1st ionisation energy decrease? | Although nuclear charge increases, it's effects are outweighed by the increase in radius and the increased shielding (to a lesser extent) |
| Across period 2 which elements experience significant falls in the ionisation energy meaning there is not a steady increase? | Beryllium to boron Nitrogen to oxygen It links to the filling of s and p orbitals |
| Why is there a decrease in 1st ionisation energy from beryllium to boron? | The start of filling the 2p sub-shell 2p sub-shell in boron has higher energy than 2s in beryllium 2p electrons in boron is easier to remove than one of the 2s electrons in beryllium 1st ionisation energy is less for boron than beryllium |
| Why is there a decrease in 1st ionisation energy between nitrogen and oxygen? | Start of electron pairing in the p orbital Paired electron in 1 2p orbital repel one another so is easier to remove an electron than in nitrogen 1st ionisation energy less for oxygen than nitrogen |
| What are metalloids/semi-metals? | Elements near the metal/non-metal divide that show inbetween properties e.g. Boron (B), Silicon (Si), Geranium (Ge), Arsenic (As) and Antimony (Sb) |
| Do all metals have the same properties? | No W (tungsten) is strong and hard, Al (Aluminium) is light and Os (Osmium) is heavy |
| How does conduction take place? | Charge must be able to move within a structure for conduction to take place |
| The structure of metal-metal bonding | Metallic bonding- Giant metallic lattice Each atom donates negative outer-shell electron which are delocalised throughout the whole structure and becomes positive ions (cations) |
| Role of cations and electrons within the metallic structure | Cations are in a fixed position maintaining the structure and shape of the metal Delocalised electrons are mobile and able to move throughout the structure Charge of electrons = charge of the cations |
| Properties of metallic bonding | Strong metallic bonds- between cations and electrons High electrical conductivity High melting and boiling points |
| Electrical conductivity in metallic bonding | Conduct electricity as solids and liquids Voltage applied to a metal, delocalised electrons can move through a structure and carry charge |
| Melting and boiling points of metallic bonding | Melting points depends on strength of metallic bonds holding the atoms in a giant metallic lattice Most metals- high temperature is needed to overcome strong electrostatic attraction between cations and electrons- have high melting and boiling points |
| Solubility of metallic bonding | Metals do not dissolve Interactions between polar solvents and charges in metallic lattice would lead to a reaction not dissolving |
| Non metal and non-metals | Covalently bonded molecules Solid state- molecules form simple molecular lattice structure held together by weak intermolecular forces |
| Which non-metals do not have simple molecular structures? | Atoms are held together by covalent bonds and form a giant covalent lattice these include: boron, carbon and silicon |
| Carbon (in it's diamond form) and silicion have a giant covalent lattice structure, but what shape structure do they form? | They are both in group 14 (4) so have 4 electrons in their outer shell so can form 4 covalent bonds It results in a tetrahedral stucture e.g. carbon in its diamond form above |
| Melting and boiling points of giant covalent lattices | High melting and boiling points Covalent bonds are strong so high temperatures needed to provide large energy quantity needed to break strong covalent bonds |
| Solubility of giant covalent lattices | Insoluble to most solvents Covalent bonds in the structure are too strong to be broken down by interactions with a solvent |
| Electrical conductivity of giant covalent lattices | Non-conductors of electricity ONLY EXCEPTION is graphene and graphite (carbon forms). In diamond (carbon) and silicon all 4 electrons are involved in covalent bonding so cannot conduct electricity. In graphene and graphite 1 electron is avaliable for conduction so can conduct electricity. |
| Graphene and graphite | Planar hexagonal layers with 120˚ bond angles due to electron pair repulsion Only 3 of the 4 electrons are used in covalent bonding, so has delocalised electrons Graphene is 1 layer of graphite Graphite has parrallel layers of heaxonally arranged carbon atoms Layers are bonded by weak London forces Both have delocalised electrons so can conduct electricity |
| Periodic trends in melting points- group 2 and 3 | |
| Why is the periodic trends in melting point across periods 2 and 3 as they are? | Melting point increases group 1-14 (4) Sharp decrease between groups 14 (4) and 15 (5)- change from giant to simple molecular structures. Also divide between metals and non-metals Melting points are compartively low from groups 15 (5) to 18 (0) Giant structures have strong forces to overcome (high melting points) and simple molecular have weak forces to overcome (low melting points) |
Want to create your own Flashcards for free with GoConqr? Learn more.