{"id":1824,"date":"2016-04-27T09:56:48","date_gmt":"2016-04-27T09:56:48","guid":{"rendered":"https:\/\/blogs.glowscotland.org.uk\/gc\/hyndsechchemu1hwrk\/?page_id=1824"},"modified":"2016-05-04T13:05:12","modified_gmt":"2016-05-04T13:05:12","slug":"review-factors-affecting-rate-of-reaction-hlo1_1b","status":"publish","type":"page","link":"https:\/\/blogs.glowscotland.org.uk\/gc\/hyndsechchemu1hwrk\/review-factors-affecting-rate-of-reaction-hlo1_1b\/","title":{"rendered":"Review [HLO1_1B] – Factors Affecting Rate of Reaction"},"content":{"rendered":"\n\n\n
\"HLO1_1B\"<\/a>
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Reaction Rate Revision Resources:<\/span><\/h3>\n\n\n\n
Written Revision Questions<\/td>\n\"HwrkRevisionIcon\"<\/a><\/td>\nMultiple Choice Revision Questions<\/td>\n\"MCQIcon\"<\/a><\/td>\n<\/td>\n<\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n

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Factors affecting the rate of Reaction<\/h3>\n

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Concentration of Reactants<\/h3>\n

(applies to reactions involving solutions only)
\nAs the concentration of reactants increases, the rate of reactions also increases.<\/span>
\nThis can be easily explained by remembering that reactions only occur when reacting particles collide. If we squeeze more reacting particles into a smaller volume (increasing the concentration), then there is a greater chance that a collision will occur. More collisions means more reactions<\/span>.
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Pressure<\/h3>\n

(applies to reactions involving gases only)
\nThe pressure of a gas is related to the number of molecules contained within a specific volume. The pressure increases when we squeeze the same number of molecules into a smaller volume.
\nAs pressure increases, the rate of a chemical reaction increases.<\/span>
\nThis observation is also easily explained by remembering that reactions only occur when reacting particles collide. If we squeeze more reacting particles into a smaller volume (increasing the pressure), then there is a greater chance that a collision will occur.
\nMore collisions means more reactions<\/span>.
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Surface Area<\/h3>\n

(applies to reactions involving solids)
\nThe greater the surface area (smaller particle size), the faster the rate of reaction.<\/span>
\nSolids can only react on their surface, the molecules under the surface are hidden, so cannot react. By breaking a larger lump into smaller pieces, a greater surface area is exposed, so there are more reacting particles, on the surface, able to take part in collisions.<\/p>\n

More collisions means more reactions<\/span>.
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Temperature<\/h3>\n

The higher the temperature, the faster the rate of reaction.<\/span><\/p>\n

Temperature measures the average<\/em> kinetic energy of the reacting particles in the substance. The graph below shows the energy of each molecule in a sample at higher and lower temperatures. The molecules have different levels of kinetic energy.<\/p>\n

According to collision theory, molecules need to collide with sufficient energy before they can react. The minimum energy required to react is called the activation energy<\/span> for the reaction.<\/p>\n

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Examining the graph, only some of the particles will have sufficient enough kinetic energy to react, and those with lower kinetic energy will not be able to react. As the temperature increases, the proportion of particles with higher energy is also increased.<\/p>\n

If more molecules have sufficient energy to react, a greater of proportion of collisions will result in a successful reaction (fruitful collisions).<\/td>\n<\/tr>\n

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Collision Geometry<\/h3>\n

For a collision to cause a reaction, moleules must collide with the correct orientation. This is best illustrated by two examples:<\/p>\n

In the reaction between ethene and hydrogen chloride, illustrated below, only if the hydrogen side of the H-Cl bond meets the carbon-carbon (double bond), will a reaction occur. If it collides with the ethene molecule in a different alignment no reaction will occur.<\/td>\n<\/tr>\n

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In this second example two hydrogen atom react to form a hydrogen molecule. No bonds are broken, only formed. In this case, because hydrogen atoms are symmmetrical in 3 dimensions collision geometry is identical regardless of where the atoms approach each other, therefore in this case, collision geometry has no impact on reaction rate – this is an unusual case.<\/td>\n<\/tr>\n
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UC Davis Chemwiki<\/a><\/td>\n<\/tr>\n

UC Davis Chemwiki<\/a><\/td>\n<\/tr>\n
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Use of a Catalyst<\/h3>\n

A catalyst reduces the activation energy of a chemical reaction. It is also said to provide an alternative pathway for the reaction to occur.\u00a0 In doing so, catalysts speed up the rate of a chemical reaction.\u00a0 More information on the action of catalysts will be found in other section of Unit 1.
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