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FlashCards por Catherine Kidd, atualizado more than 1 year ago
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Criado por Catherine Kidd
quase 10 anos atrás
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| Questão | Responda |
| Define the term first ionisation energy Define the term successive ionisation energy | First ionisation energy - energy required to remove one mole of electrons from one mole of atoms in the gaseous state to form one mole of gaseous 1+ ions Successive ionisation energy - energy required to remove each electron in turn. |
| Explain the factors that influence ionisation energy | The ionisation energy must be sufficient to overcome the attraction of the negative electrons to the positive nucleus. Outer electrons are lost first as they are furthest from the nucleus and so experience the least nuclear attraction. The strength of this attraction is influenced by 3 factors 1. Atomic radius - the greater the atomic radius, the smaller the nuclear attraction and therefore the smaller the ionisation energy, due to increased distance from the nucleus and increased shielding 2. Nuclear charge - the more protons in the nucleus the greater the positive charge of the nucleus and the greater the nuclear attraction, therefore the greater the ionisation energy. 3. Electron shielding - the repulsion between electrons in different shells, the more inner shells the greater the shielding effect, which overcomes the nuclear attraction from the nucleus, therefore decreasing the ionisation energy. |
| Predict the number of electrons in each shell of an atom from its successive ionisation energies Predict the group of an element from its successive ionisation energies. | As each electron is removed, there is less repulsion between electrons causing each shell to be drawn slightly closer to the nucleus, decreasing the atomic raduis of the atom. This increases the nuclear attraction felt by the electrons causing the ionisation energy to increase. Each successive ionisation energy is larger that the one before. There is a large increase as the shell from which electrons are being removed changes, allowing one to predict the number of electrons in each shell. The group of the element can be guessed by looking counting ionisation eneries before the first large jump to see how many electrons are in the outermost shell. |
| Define an orbital Describe the shapes of s and p orbitals | Orbital - a region in space that can hold up to two electrons with opposite spins. An s orbital is spherical in shape A p orbital has a 3D dumb-bell shape |
| State the numbers of orbitals making up the s-, p- and d- subshells State the number of electrons that occupy s-, p- and d- subshells | Each electron shell is made up of orbitals with the same principal quantum numer (energy level), orbitals of the same type within shells are grouped into sub-shells. s- subsell - one s- orbital = 2 electrons p- subshell - three p- orbitals = 6 electrons d- subshell - five d- orbitals = 10 electrons f- subshell - seven f- orbitals = 14 electrons |
| State the number of electrons that can fill the first three shells | The first shell contains only one s- orbital/subshell so can hold 2 electrons The seconds shell contains one s- subshell and one p- subshell so can hold 8 electrons The third shell contains one s- subshell, one p-subshell, and one d- subshell so can hold 18 electrons The fourth shell contains one s- subshell, one p- subshell, one d- subshell and one f- subshell so can hold 32 electrons |
| Describe the relative energies of s-, p- and d- orbitals in the first 3 shells | The subshells within a shell have slightly different energy levels, despite being grouped. Within a shell, the energy of each subshell increases in the order s, p, d, and f. Lowest available energy levels must be filled first and each energy level must be filled before the next can start to fill. All orbitals within a subshell also fill singly before pairing. There is an irregularity between the third and fourth shells whereby the s subshell of the fourth shell has a lower energy level than the d subshell of the third shell. Therefore the fourth shell must begin to fill before the third shell is completely filled. |
| Deduce electron configurations of atoms, and ions given their ionic charge | Electron configuration is the arrangement of electrons in an atom. Orbitals fill from the lowest energy level up so atoms with fewer electrons will have only the lowest energy levels filled. The shorthand used to to show electron configurations takes the form nx^y where: n shows the shell x the type of orbital ^y the number of electrons in the subshell. e.g N = 1s^2 2s^2 2p^3 The subscript numbers should add up to the total number of electrons that an atom of the element contains, here 7. For ions, the process is the same, but the charge must be taken into account, therefore we take the electronic configuration of an atom of the element and the take away electrons (positive ions) from the highest filled/partially filled energy level or add electrons (negative ions) to the lowest available energy level. However, again the exception is with the 4s- subshell which, once filled, becomes of slightly higher energy than the 3d- subshell, therefore although the 4s- subshell is filled first, electrons are also removed from it first. |
| Deduce electron configurations of atoms, and ions given their ionic charge | Electron configurations may be abbreviated by showing the noble gas configurations. Here, the inner shell configurations are shown by stating the name of the closest previous noble gas e.g [He] and then only showing the electron configuration of the outermost shell e.g. Li = [He]2s^2 |
| Classify the elements into s, p and d blocks. | The block of the periodic table in which an element sits demonstrates the type of orbital in which its outermost electron is held. The row then shows the shell in which the orbital is (i.e big number) and the column shows how many electrons occupy the outermost subshell (i.e small/subscript number |
| Define the term ionic bonding Describe the term covalent bond | Ionic bonding - the electrostatic attraction between oppositely charged ions. (all ionically bonded ions form giant ionic lattices whereby each ion is surrounded by oppositely charged ions - they are non-molecular) Covalent bond - a shared pair of electrons |
| State the formulae for common ions | Nitrate ion: NO3^ - Carbonate ion: CO3^2- Sulphate ion: SO4^2- Ammonium ion: NH4^+ |
| Describe dative covalent bonding | In a dative covalent bond, one of the atoms supplies both of the electrons involved in the covalent bond e.g the ammonium ion (NH4) contains three single covalent bonds between the nitrogen atom and a hydrogen atom, and one dative covalent bond where the nitrogen provides both of the electrons in bonding with a H+ ION. (i.e proton) e.g The oxonium ion (H30+) where oxygen provides both electrons in bonding with one of the hydrogens (ion). |
| Describe covalent bonds that disobey octet rule. | 1. Not reaching the octet Be and B (in period 2) do not have enough unpaired electrons to form a noble gas configuration in a compound. e.g Boron trifluoride, BF3. B has each of its 3 unpaired electrons paired in a covalent bond, while all three fluorine atoms reach an octet. 6 electrons in B outer shell. |
| Describe covalent bonds that disobey the octet rule | 2. Expanded octet From period 3, more outer-shell electrons are able to take part in bonding, such that one of the bonding atoms in the resulting molecule may end up with more than eight electrons in its outer shell. Non-metal group 5 either form 3 or 5 covalent bonds Non-metal group 6 eother form 2,4, or 6 covalent bonds Non-metal group 7 either form 1,3,5, or 7 covalent bonds e.g sulfur hexafluoride, SF6 Sulfur forms 6 covalent bonds, pairing each of its unpaired electrons so 12 electrons occupy its outer shell. All 6 fluoride atoms obey the octet rule. The maximum number of electrons that can pair up always pair up. |
| Explain how electron repulsion theory can determine the shape of a molecule | The shape of a molecule or ion is determined by the number of electron pairs in the outer shell surrounding the CENTRAL atom As all the electrons have negative charge, the electron pairs repel each other as far away as possible. |
| Explain the shapes of, and bond angles in, molecules and ions with up to 6 electron pairs | A molecule with three bonding pairs results in a 2D trigonal planar shape with a bond angle of 120 degrees. |
| Explain the shapes of, and bond angles in, molecules and ions with up to 6 electron pairs | A molecule with four bonding pairs has a bond angle of 109.5 degrees, forming a tetrahedral shape |
| Explain the shapes of, and bond angles in, molecules and ions with up to 6 electron pairs | A molecule with 6 bonding pairs (e.g SF6) forms a octahedral shape, with a 90 degree bond angle |
| Explain the shapes of, and bond angles in, molecules and ions with up to 6 electron pairs | Lone pairs repel more than bond angles as they have a slightly higher electron density, and therefore decrease the bond angles of the molecule by 2.5 degrees per lone pair. Therefore, as a molecule with four bonding pairs (tetrahedral) has bond angles of 109.5 degrees, the bond angles of a molecule with 3 bonding pairs and one lone pair will be 2.5 degrees less than that, at 107 degrees. This type of molecule is pyramidal in shape. |
| Explain the shapes of, and bond angles in, molecules and ions with up to 6 electron pairs | A molecule with 2 bonding pairs and 2 lone pairs will have a bond angle of 104.5 degrees (5 degree less than tetrahedral) giving it a nonlinear shape or angular line shape |
| Explain the shapes of, and bond angles in, molecules and ions with up to 6 bonding pairs. | Double bonds are treated as bonding regions when determining the shape of a molecule, and behave the same way/have the same strength as bonding pairs. Therefore a molecule with 2 bonding regions, which will repel each other as far away as possible, has a bond angle of 180 degrees and a linear shape. |
| Explain the shapes of, and bond angles in, molecules and ions with up to 6 bonding pairs. | The same principles can be applied to ions as were applied to molecules that is, -electrons repel as far away as possible - lone pairs repel 2.5 degrees more than bonding pairs (take a tetrahedral shape with 107.5 degrees as a starting point) - double bonds are treated as bonding regions and act as bonding pairs. |
| Define the term electronegativity | Electronegativity - a measure of the ability of an atom to attract electrons in a covalently bonded pair towards it. |
| Explain how a permanent dipole may arise | A non-polar bond will form when the two atoms involved in the covalent bond are of very similar electronegativities resulting in an even distribution of electrons between them - neither atom attracts the electron pair more strongly than the other. e.g H-H A polar bond will form when the two atoms involved in the covalent bond are of different electronegativities resulting in an uneven distribution of electrons as the more electronegative atom will attract the bonding pair more strongly. e.g H-Cl (Cl being the more electronegative) This uneven distribution creates a small charge difference across the bond, a permanent dipole. Greater electronegativity difference = greater permanent dipole. |
| Explain how a permanent dipole may arise | A molecule with polar bonds is not necessarily polar overall. A polar molecule has an overall dipole when you take into account any dipoles across the bonds. In molecules which are symmetrical, the dipoles of the individual bonds cancel out resulting in a non-polar molecule. e.g CF4 Molecules with polar bonds that are asymmetrical (have a positive and a negative end/side) are polar molecules e.g H20 |
| Describe intermolecular forces based on permanent dipoles | Intermolecular forces are very weak in comparison to ionic and covalent bonds, they are caused by small attractive forces between oppositely charged dipoles. Permanent dipole-dipole forces form when the permanent dipole of one molecule attracts the permanent dipole of another. |
| Describe intermolecular forces based on induced dipoles | Van der Waals forces exist between all molecules, polar and non-polar, and are formed from very weak attractions between instantaneous dipoles in neighbouring molecules. Instantaneous dipoles result from the movement of electrons within shells. At any moment there may be an uneven distribution of electrons within the shells resulting in a small charge across the molecule, an instantaneous dipole. This instantaneous dipoles induces a dipole in neighbouring molecules, which in turn induce dipoles in their neighbouring molecules. The attractive forces between these dipoles are Van der Waals forces. Van der Waals forces are weaker that dipole-dipole forces, but increase with an increasing number of electrons. The more electrons the bigger the induced dipole and therefore the greater the attractive forces between the molecules. (and therefore the higher the boiling point - so boiling point increases as you move down the group) |
| Describe hydrogen bonding, including the role of a lone pair, between molecules containing -OH and -NH groups | Molecules containing O-H and N-H bonds are polar with particularly strong dipoles, as O and N are particularly electronegative. The permanent dipole-dipole interaction between these molecules is a hydrogen bond. This occurs as the delta positive hydrogen atom (electron deficient as O or N strongly attracts the bonding pair towards them) attracts the lone pair of electrons on the delta negative O or N on a different molecule. e.g in H2O |
| Describe and explain the anomalous properties of H2O resulting from hydrogen bonding | Ice is less dense than water because the hydrogen bonds in the open lattice of ice holds the water molecules apart from each other. However, when it melts the rigid hydrogen bonds collapse allowing the water molecules to move closer together. Water has high melting and boiling points because of the stronger hydrogen bonds that add to the weaker Van der Waals forces. The hydrogen bonds must be overcome in order to separate the molecules. Hydrogen bonds also give water a higher surface tension and viscosity. |
| Define metallic bonding. Describe structures as giant metallic lattices. | Metallic bonding is the electrostatic attraction between positive metal ions and delocalised electrons. A giant metallic lattice is a 3D structure of positive ions and delocalised electrons bonded by strong metallic bonds. Metallic lattices have: 1. High melting and boiling points - there is a strong electrostatic attraction between the positive ions and the negative electrons so high temperatures are needed to overcome the bonds and dislodge the fixed ions 2. Good electrical conductivity - delocalised electrons are able to move freely throughout the lattice and carry the charge 3.High malleability and ductility - delocalised electrons are free to move throughout the structure, allowing the layers of ions to slide over/past each other. Alloys are mixtures of metals which are harder than a pure metal because the layers become distorted such that they are no longer able to slide past each other due to the differently sized ions of the metals. |
| Describe structures as giant ionic lattices, with strong ionic bonding. Describe the properties of giant ionic lattices | All ionic compounds exist as giant ionic lattices in a solid state. Each ion attracts oppositely charged ions surrounding it, forming the giant ionic lattice. |
| Describe the properties of giant ionic lattices | 1. High melting and boiling points - the electrostatic attraction between the oppositely charged ions is very strong therefore a lot of heat energy is required to overcome the forces and separate the ions. The greater the charge on the ions, the stronger the electrostatic attraction and the higher the boiling/melting point. 2. High electrical conductivity when molten/dissolved in water - as a solid, ions are in fixed positions but when in liquid state the lattice breaks down, allowing the ions to move freely and carry the charge. 3. High solubility in polar solvents e.g water - water has polar molecules which are able to break down the giant ionic lattice by surrounding each ion to form a solution. The delta positive H surrounds the negative ions and the delta negative O surrounds the positive ions. |
| (Example of giant ionic lattice solubility in apolar solvent) | e.g NaCl |
| Describe structures as simple molecular lattices Describe the properties of simple molecular lattices | Simple molecular lattice - 3D structure of molecules bonded together by weak intermolecular forces. However, the atoms within each molecule are held together by strong covalent bonds. Simple molecular lattices have: 1. Low melting and boiling points - in order for the compound to change state, only the weak Van der Waals forces need to be broken, therefore little heat energy is required to overcome them. 2. No electrical conductivity - do not contain charged particles. 3. High solubility - dissolve in non-polar solvents because Van der Waals forces form between molecules of the simple molecular lattice and non-polar solvent, which weaken the simple lattice structure. |
| Describe structures as giant covalent lattices Describe the properties of giant covalent lattices | Giant covalent lattice - 3D structure of atoms bonded together by stong covalent bonds e.g diamond, graphite, silicon dioxide. Giant covalent structures have: 1. High melting and boiling points - the strong covalent bonds need to be broken in order for the compound to change state, which requires a lot of heat energy. 2. No electrical conductivity - there are no free charged particles to carry the charge. Graphite is an exception as each carbon atom is only bonded to three other carbons atoms therefore one outer shell electron from each atom is not used in bonding, resulting in a sea of delcoalised electrons which can more throughout the structure and carry the charge. 3. No solubility - the covalent bonds are too strong to be broken by either polar or non-polar solvents. 4. Diamond is hard - tetrahedral shape allows external forces to be distributed throughout the lattice Graphite is soft - hexagonal layer structure with only Weak Van der Waals forces acting between the layers, allowing them to easily slide over each other. |
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