2118512
Description
Flashcards by James Farley, updated more than 1 year ago
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Created by James Farley
over 9 years ago
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| Question | Answer |
| Atomic number | The number of protons in the nucleus of an atom. |
| Mass number | The number of protons and neutrons in the nucleus of an atom. |
| Isotopes | Atoms of the same element with different numbers of neutrons. |
| Why do isotopes have similar chemical properties? | They have the same number of electrons in their outer shell. |
| Relative isotopic mass | The mass of an isotope compared to 1/12th of the mass of one atom of carbon 12. |
| Relative atomic mass | The weighted mean mass of an atom compared to 1/12th of one atom of carbon 12. |
| Protons | Relative mass: 1 Relative charge: +1 Position in the atom: Nucleus |
| Neutrons | Relative mass: 1 Relative charge: 0 Position in the atom: Nucleus |
| Electrons | Relative mass: 1/2000 Relative charge: -1 Position in the atom: Shells |
| Relative Molecular mass | The weighted mean mass of a molecule relative to 1/12 of one atom of carbon 12. |
| Orbital | A volume of space that can hold up to two electrons with opposite spins. |
| Order of the orbitals | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 (max 36 electrons) |
| First ionisation energy | The amount of energy required to remove one electron from each atom in one mole of gaseous atoms (to form one mole of positive gaseous ions.) |
| Successive ionisation energies | Second ionisation energy: O+ (g) ----> 0 2+ (g) + e- Third ionisation energy: O 2+ (g) ----> O 3+ (g) + e- |
| Trend in successive ionisation energies | Removal of additional electrons results in an increase in ionisation energy because: Same number of protons are now attracting fewer electrons in the ion One less electron means there is less electron-electron repulsion The O+ ion is smaller than the O atom, therefore the next electron to be removed will be closer to the nucleus Next electron will experience greater nuclear attraction. |
| CARS: across a period | Increase in ionisation energy: C inc - less electrons added to same no. of protons A inc R dec - extra electron shell added S eq - electrons added to same shell |
| CARS - down a group | Ionisation energy decreases: C inc - outweighed by radius and shielding A dec R inc - extra electron shell added S - more electron-electron repulsion |
| Acid + Metal ---> 2HCl (aq) + Mg (s) ---> | Salt + hydrogen gas MgCl2 (aq) + H2 (g) |
| Acid + Metal Oxide ---> 2HCl (aq) + CaO (s) ---> | Salt + Water CaCl2 (aq) + H20 (l) |
| Acid + Metal hydroxide ---> HNO3 (aq) + NaOH (s) ---> | Salt + Water NaNO3 (aq) + H20 (l) |
| Acid + Metal carbonate ---> H2SO4 (aq) + CaCO3 (s) ---> | Salt + Water + Carbon dioxide CaSO4 (aq) + H2O (l) + CO2 (g) |
| Ionic bond | The electrostatic force of attraction between oppositely charged ions (occurs between a metal and a non metal.) |
| Covalent bond | A shared pair of electrons (occurs between two non metals.) |
| Dative covalent bond | A shared pair of electrons, both of which come from the same atom. E.g. NH4+ formed by reaction of NH3 + H+. |
| Metallic bond | The electrostatic force of attraction between positive metal ions and delocalised electrons. |
| Exceptions of the octet rule | When there are not enough electrons to complete the octet. E.g. Be & B in period 2 when they are the central atom. |
| Expansion of the octet rule | When each atom has more than 8 electrons. E.g. Groups 5-7 from period 3 downwards when they are the central atom. |
| Linear shaped molecules | E.g. Becl2 BeH2 CO2 2 bonded pairs 0 lone pairs 180° bond angles |
| Trigonal planar shaped molecules | E.g. BF3 BCl3 BBr3 SO3 3 bonded pairs 0 lone pairs 120° bond angles |
| Tetrahedral shaped molecules | E.g. CH4 SiH4 NH4+ CCl4 4 bonded pairs 0 lone pairs 109.5° bond angles |
| Octahedral shaped molecules | E.g. SF6 SeF6 6 bonded pairs 0 lone pairs 90° bond angles |
| Pyramidal shaped molecules | E.g. NH3 NCl3 PCl3 3 bonded pairs 1 lone pair 107° bond angles (-2 . 5° for each lone pair) |
| Non-linear shaped molecules | E.g. H2O H2S 2 bonded pairs 2 lone pairs 104.5° bond angles |
| Electronegativity | The ability of an atom to attract a pair of electrons in a covalent bond towards itself. |
| Pauling scale | Electronegativity increases to the top right of the periodic table, with F being the most electronegative. |
| Permanent dipole | A difference in electronegativity between 2 atoms results in a small charge difference. Eg. HCl H △+ Cl△- electrons always closer to Cl atom. |
| Polar bonds in a non polar molecule | A molecule may have polar covalent bonds because the molecules are symmetrical. The dipoles act in opposite directions and cancel each other out. |
| Permanent dipole-dipole forces | One molecule can attract the opposite permanent dipole in a neighbouring molecule. |
| Van der waals' forces | Forces arise due to movement of electrons (uneven distribution.) This causes an instantaneous dipole (temporarily.) This temporary dipole induces a dipole in neighbouring molecule. Dipoles attract one another (weak VDW's) |
| Increasing van der Waals' forces | VDW's forces increase with an increase in the number of electrons. More electrons result in larger temporary and induced diploles, which results in a greater force of attraction between molecules. Therefore boiling point increases. |
| Hydrogen bonding | Special permanent dipole-dipole force between O-H or N-H bonds. The H△+ attracts the lone pair of electrons in the O△-/N△- in neighbouring molecule. |
| Special properties of water | Ice is less dense than water: the hydrogen bonds hold H2O molecules further apart in an open lattice. H2O has a higher than expected boiling point: hydrogen bonds are the strongest intermolecular force therefore require more energy to break. |
| Simple molecular lattice | E.g. H2 I2 H2O Low melting/boiling points: weak VDW's forces require little energy to be broken. Don't conduct electricity: no free charged particles. |
| Giant covalent lattice | E.g. Si, SiO2, diamond&graphite High melting/boiling point: strong covalent bonds require lots of energy to break. Don't conduct electricity: no free moving charged particles |
| Diamond | Giant covalent lattice. Doesn't conduct electricity: no free moving charged particles. Hard: strong covalent bonds throughout tetrahedral structure. |
| Graphite | Giant covalent lattice with layers of carbon atoms. Conducts electricity: delocalised electrons between carbon layers. Soft: weak VDW's forces between the layers allowing them to slide. |
| Ionic compounds | High melting/boiling point: strong electrostatic force of attraction between oppositely charged needs to be broken. Don't conduct in solid state: fixed position. Conduct when liquid/aqueous: ions free to move. Soluble in polar solvents, surround and break down lattice. |
| Metal bonding | The strong electrostatic force of attraction between positive metal ions and delocalised electrons. High melting/boiling point: strong electrostatic force of attraction. Conduct electricity: delocalised electrons are free to move. Ductile: drawn out (wires.) Malleable: hammered into shape. |
| ionic charge affecting melting point of a metal | E.g. Na+ Mg2+ Al3+ As ionic charge increases the number of delocalised electrons increases. Therefore attraction increases between delocalised electrons and metal ions (stronger metallic bonding.) |
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