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AS physical Defs
All physical definitions from AQA AS chemistry year 12
| Term | Definition |
|---|---|
| Mass number (A) | Total number of protons & neutrons present in the nucleus of an atom |
| Atomic number (Z) | Total number of protons present in the nucleus of an atom |
| Mass spectrometry | Analytic technique that determined the molecular mass of a compound. Involves ionisation, separation, & detection of an ion based on their mass to charge (m/z) ratio & gives relative abundance of different isotopes |
| Two types of ionisation | Electron impact ionisation - electrons are fired from electron gun to knock of an electron from outer shell Electrospray ionisation - passes it by a positive terminal to add an H+ ion |
| Process of mass spectrometry | Ionisation (compound is given positive charge) acceleration (negatively charged plate gives ions same kinetic energy but different momentum), detection (ions reach detection plate, receive electron, & give off current proportional to their m/z value |
| Equation for ionisation is mass spectrometry | Electron impact = X(g) — X+(g) + e- Electrospray = X(g) + H+(g) — XH+(g) |
| Relative atomic mass (Ar) | Ratio of weighted average mass of all isotopes of an element (in terms of relative abundance) compared to 1/12 of the mass of carbon-22 |
| Relative molecular mass (Mr) | The sum of the average mass of all atoms included compared to 1/12 of the mass of carbon-12 |
| Avogadro constant | The approximate number of particles in 1 mole of substance, 6.022 x 10^23 |
| Volume of a mole of gas | Under standard temperature & pressure, 1 mole of gas occupies 22.5dm^3 |
| Ideal gas equation | pV = nRT P = pressure/Pa V = volume of container/m^3 n = number of moles R = gas constant = 8.31 Jmol^-1K^-1 |
| What the ideal gas equation assumes | Particles don’t collide or interact by attractive or repulsive forces, individual particles occupy no volume, they move randomly & independently from each other, they don’t lose kinetic energy from colliding with container walls (i.e, perfectly elastic) |
| How to calculate empirical formula | Calculate mass of each element in a 100g sample of the compound, divide each element mass by their Mr. Divide all number of moles by smallest number of moles to find molecular ratio & multiply all numbers until there are only whole numbers |
| Atom economy | The theoretical proportion of reactant atoms that become part of a desired product in the chemical reactions. Calculated by dividing molar mass of desired product(s) by molar mass of all reactants and x100 |
| Yield | Amount of desired product obtained from a chemical reaction |
| Ionic equation | Chemical equation representing transfer of electrons between species to create ionic compounds. Most reactants & products are written as ionic forms & have state symbols that are nearly always aqueous (can sometimes be a gas) |
| Half equation | A chemical reaction representing either the oxidation or reduction of a redox reaction focusing on transfer of electrons between species to make a product. 2 can be put together to make an ionic equation |
| Ionic bonding | Type of bonding involving electrostatic attraction between oppositely charged ions in a lattice. Ions can bond to many oppositely charged ions |
| Covalent bonding | Strong electrostatic attraction between a bonding pair of electrons (at least 1 from each atom) & the positively charged nuclei of the 2 bonded atoms. Commonly forms between non-metals |
| Dative bonding | Type of covalent bond between 2 atoms where the shared electron pair originates from only 1 atom, commonly represented by an arrow |
| Metallic bonding | Electrostatic attraction between positively charged metal ions & ‘sea’ of delocalised electrons. This happens when metal atoms dissociate from their valence electrons which can now move freely through the metallic lattice |
| Physical properties of ionic compounds | High melting point (strong electrostatic forces), brittle crystalline structure (when pressure’s applied, same charged ions repel each other & break lattice), conducts electricity when melted or in solution due to being charged |
| Physical properties of metallic structures | High thermal & electrical conductivity (many delocalised electrons), malleability & ductility (non-directional bonds, sliding layers of cations without breaking structure), high melting point (strong attraction), can still bond in liquid form |
| Physical properties of covalently bonded molecular compounds | Variable melting & boiling points due to changes in molecular shape, size, & strength of intermolecular forces. Poor electrical conductivity due to localised electrons |
| Physical property of covalent macromolecular compounds | Strength & rigidity from high number of bonds & interactions between chains. High melting & boiling point since more binds take more energy to overcome. Can be soluble if is polar in nature, extremely low conductivity due to localised electrons |
| Bonding pairs | Pair of electrons shared between 2 atoms in a chemical bond |
| Lone pairs | Pair of electrons in valence shell of an atom not involved in bonding |
| Polar bond | 2 atoms of very different electronegativities bond covalently causing uneven distribution in e- density cloud making it gather to the more electronegative element making it partially negative while the other is partially positive |
| Dipole moment | Measure of bond polarity or separation of charges in a molecule |
| Permanent dipole-dipole forces | Permanent significant uneven distribution of e- density due to a polar bond. These partial charges can be attracted to opposite charged poles in other polar molecules |
| Induced dipole-dipole forces (Van Der Waals) | Weak intermolecular forces between non-polar molecules. Random temporary fluctuations in e- density affect e- density in nearby molecules creating temporary poles which can attract each other |
| Hydrogen bonding | Strongest intermolecular force, occurs between extremely polar molecules formed when hydrogen bonds to a very electronegative element (nitrogen, oxygen, fluorine) & requires a lot of energy to overcome |
| How intermolecular attraction affects melting & boiling points | Small molecules with VDW generally have low, but bigger molecules have more e- density so it’s high for them. D-D interactions normally have higher melting than VDW but depends of sizes of atoms. Hydrogen & ionic bonding are nearly always the highest |
| Enthalpy change (+standard enthalpy change) | Heat change of chemical reactions where reactants are converted into products. Standard is the change under standard conditions with all chemicals in their standard states |
| Enthalpy of combustion | Amount of heat energy released when 1 mole of substance reacts with excess oxygen under standard conditions |
| Enthalpy of formation | Energy change when 1 mole of product is formed from its constituent parts under standard conditions |
| Formula for calculating enthalpy change | Q = mcθT Q = enthalpy change m = mass of substance heated (normally water) c = specific heat capacity (4.18 for water) θT = change in temperature |
| How to turn enthalpy change into molar enthalpy | (Enthalpy change/1000)/number of moles |
| Bond enthalpy | Overall energy change when 1 gaseous mole of a particular bond is broken under standard conditions |
| Mean bond enthalpy (+ explanation) | Bond enthalpy of a particular bond averaged over a range of compounds. This is because each bond is broken in a different condition |
| Kinetics | Rate of chemical reactions and factors that influence it |
| Rate of reaction | How quickly concentration of reactant decreases (or concentration of product increases) per unit of time |
| Activation energy | Minimum amount of energy for a chemical reaction to occur |
| Collision theory | Chemical reactions only occur when particles collide with each other in the correct orientation with sufficient energy |
| Why most collisions don’t lead to reactions | Incorrect orientation, inadequate kinetic energy, insufficient collision frequency, or lack of catalyst |
| Maxwell-Boltzmann distribution curve | Graph displaying molecular velocities of a sample of gaseous molecules at a given temperature. X axis represents speed in kinetic energy, Y axis represents frequency in number of particles having a specific speed |
| Where Maxwell-Boltzmann curves originate | At the origin point because 0 particles have 0 energy |
| Why might a small increase in energy lead to a large increase in rate of reaction (exponential rather than linear) | Particles must collide with sufficient energy to react, increasing the temperature increases proportion of particles with kinetic energy </=activation energy causing more successful collisions per unit of time |
| Effect of concentration changes on collision frequency | If concentration of particles increases then more particles per set volume so higher probability of particles with activation energy colliding which increases rate of reaction & vice versa |
| Effect of pressure changes on collision frequency | As pressure increases, available volume decreases causing gas particle speed to increase. As average speed increases, successful collisions increase & vice versa |
| Catalyst | A substance that lowers the minimum activation energy of a reaction without being changed in composition or amount by providing an alternative reaction route that requires lower activation energy. Affects the forwards and backwards reactions equally |
| Chemical/dynamic equilibrium | Dynamic state in reversible reactions where forward & backward reactions occur at same rates. Concentrations of products & reactants remain constant despite the fact the reactions are still occurring |
| Forward reaction | Reactants undergo chemical changes to become products, represented by reactants on left hand side of reaction and products on the right |
| Backward reaction | Products formed in a chemical reaction either react with each other or decompose to form their original reactants |
| Reversible reactions | Chemical reaction where both forwards and backwards reactions are simultaneously feasible (though often to different extents) |
| Le Chatellier’s principle | If an external change is applied to a system at equilibrium, the system will adjust itself to oppose the change and restore a new equilibrium |
| Effects of temperature change on equilibrium | One direction will be exothermic while the other will be endothermic (not based on which direction); if temperature increases, equilibrium will shift in the endothermic direction & vice versa for the exothermic direction |
| Effect of pressure change on equilibrium | The direction that leads to the fewest gaseous molecules will be favoured if pressure increases regardless of concentration or mass of chemicals on either side & vice versa |
| Effect of concentration change on equilibrium | If reactant concentration increases, forward direction is favoured & equilibrium is shifted right. Vice versa |
| Equilibrium constant (Kc) | Ratio of concentration of products to reactants under standard conditions for a reversible reaction. Commonly represented as aA + bB -/- cC + dD. Lowercase are concentrations & uppercase are number of molecules. Also written as Kc = [C]^c[D]^d/[A]^a[B]^b |
| Homogenous system | Mixture where components are uniformly distributed throughout the mixture making the entire composition & properties the same at any point in the mixture |
| Oxidation | A atom, ion, or molecule loses an electron. Commonly accompanied by the gain of oxygen &/or loss of hydrogen |
| Oxidising agent | Substances that cause oxidation by accepting electrons from other species in a chemical reaction, thereby being reduced themselves |
| Reduction | An atom, ion, or molecule gains an electron. Commonly accompanied by the loss of oxygen &/or gain of hydrogen |
| Reducing agent | Substances that cause reduction by donating electrons to other species in a chemical reaction, thereby being oxidised themselves |
| Redox | Chemical reaction involving transfer of electrons between different chemical species resulting in one or more being oxidised & one or more being reduced |
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