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Chemistry 1050 lecture notes summerized week 1 and 2.pdf

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CHEM 1050
Harvey Marmurek

Chemistry 1050 Lecture notes summarized WEEK 1 and 2 Notes Intro to thermodynamics:  It is the interconversion of different forms of energy  Examples: o Heat, mechanical work, electrical work, chemical energy  It is an essential function for understanding the physical and biological sciences  It is basis for the physical description of living matter Energy Transfer: Chemical Reactions: - They involve energy transfer between: the system and the surroundings The system is defined as the: - The part of the universe under observation - Examples include: engine, electrical cell, reaction in a flask The surroundings are defined as: - The rest of the universe (everything outside of the system) - Therefore the universe is made up of the system and the surroundings The SYSTEM can be: OPEN: - Both mass and energy may enter and leave - Examples: beakers, open water bottles CLOSED:  Energy can be exchanged but no mass can enter or leave  Examples: sealed container, closed water bottle ISOLATED:  Nothing enters or leaves  Example: thermos 1 State Function:  The state of the system can be defined by a set of variables  Examples: T, P, V, n, etc. o Temperature does not change with size  Only initial and final points effect this function  The path is not important  2 classifications o Intensive  Independent of size  Example: T, P, etc. o Extensive  Dependent upon size  Example: mass, V, n, etc. Energy Transfer: By transferring energy into or out of the system, the internal energy of the system changes  The internal energy (U) of a system is defined as the sum of the kinetic and potential energy  U can change in 2 ways: 1. Heat (q) 2. Work (w)  Work is mechanically based  Heat into the system, thermal surroundings lose heat; system gain heat  U cannot be measured directly we can only measure changes in delta U  D U reaction = U products – U reactants  By any combination of q and w, we can change U:  THIS IS THE FIRST LAW OF THERMODYNAMICS!!!!!! Energy cannot be created or destroyed, only transformed from one form to another  The energy can also be transferred into or out of the system  This is also the law of the conservation of energy Heat:  Heat content is defined by temperature, T  If 2 objects are the same temp no heat transfer will occur  Heat flows from high to low  The cause of energy flow is molecular motion Heat and Chemical Reactions:  The amount of heat depends on the size of the system, this makes q an extensive property  Heat given off or absorbed in a constant pressure system is known as the Enthalpy Change (Delta H) 2  The Enthalpy Change, Delta H, for a process can be defined as either being endothermic or exothermic ENDOTHERMIC:  Heat absorbed by the system from the surroundings  Positive EXOTHERMIC:  Heat absorbed from the system to the surroundings Enthalpy: 1. Exothermic reactions ( Delta H is less than 0)  Heat is given OFF to the surroundings  This arises due to the differences in bonding energy of reactants and products  ALL combustion reaction are exothermic 2. Endothermic reactions (Delta H is greater than 0)  Heat is absorbed by the reaction What is Enthalpy? - It is a measure of the difference in heat content, H, between the products and the reactants of a reaction Heat vs. Temperature Heat: depends on mass and speed (quantity and speed) Temperature: speed of molecules (speed only) The measure of Heat: Calorimetry CALORIMETRY: is a technique that has been developed to measure the heat change of a system “ Calorimetry” is derived from greek and latin, calor, meaning heat, and metry, meaning to measure Need to measure temperature change (change in heat) Calorimetry: Experimentally, we need to:  Measure the temperature change during a reaction  Convert this change in temperature to a change in energy  To do this, we use HEAT CAPACITY (C), which is defined as the energy required to raise the temperature of a mass or substance by 1 degree celeci
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