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# CHEM 1020 Module 1 - Chemical Energetics

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School
Department
Chemistry
Course
CHEM1100
Professor
Gwen Lawrie
Semester
Spring

Description
CHEMICAL ENERGETICS KEY VARIABLES IUPAC standard pressure = 273.15 k and 1 bar ABSOLUTE TEMPERATURE Measured in Kelvin 0K = -273.15 C (absolute zero) All calculations done in Kelvin must not be negative values (lowest temperature is 0K) GAS PRESSURE Particles in constant motion collide with container walls (surroundings) – each have a mass and velocity Pressure = measure of collisions per unit area per unit time Measurement of gas pressure SI units: pascal (Nm ) Non-SI units:  1 atm (101.3kPa or 760 mm Hg)  1 bar = 100 kPa  1 torr = 1/760 atm Barometer: pressure of the atmosphere on the surface of the mercury balances the pressure of the column of mercury Contemporary devices: pressure sensors (transducers) measure the difference between two pressure points electronically Pressure of a gas changes with:  Volume  Temperature  Number of moles Boyle’s Law V1/V2 = P2/P1 (at constant temperature) Pressure exerted by a gas held at a constant temperature is inverse to the volume of the gas Volume halved = pressure doubled * Gas comprised of Gas molecules hit loosely spaced walls of container molecules moving at more frequently and randm exert increas pressure Gas compressed, pushing the molecules occupy less volumeo * Applies only to ideal gases Charles’ Law V1/T1 = V2/T2 (at constant pressure) If the pressure of a gas remains constant, the volume of the gas will increase as the temperature increases Increased temperature = gas takes up more space Avogadro’s Law Gas volume doubles as the amount of gas doubles (at constant temperature and pressure) IDEAL GASES Rapid motion of gas particles + collisions with the walls of their container = pressure PROPERTIES OF IDEAL GASES Ideal gas theory: pV = nRT P = pressure (often in Pa) V = volume (often in m ) N = molar amount of gas R = gas constant (8.314 J mol K ) T = temperature (always in K)  No fixed volume or space – will expand to occupy all available space  Move independently and randomly of each other  Exert a pressure  No loss of kinetic energy during collisions  No intermolecular forces between molecules KINETIC THEORY Gas molecules have mass (m) and velocity (u) – therefore have kinetic energy: Molecules are in constant random motion – energy of molecule is related to speed Molecules move faster as temperature increases Smaller molecules tend to move faster than larger molecules At a given temperature all gases have the same molecular kinetic energy distribution Average and total kinetic energy For a gas at a given pressure and temperature, the average energy is independent of the amount of gas To find average kinetic energy: Where T = temperature (K)A N is Avogadro constant and R is the gas constant (8.314J mol K ) To find the total kinetic energy of one mole, multiply by Avogadro constant: 1 mole of gas occupies 22.7L at STP (0 C and 1 bar) 1 L = 1 x 10 m -3 DALTON’S LAW OF PARTIAL PRESSURES P T P +1P +2... + P I The total pressure of a mixture of gases is the sum of the pressures that each gas would exert if it were alone in the container REAL GASES PROPERTIES OF REAL GASES At sufficiently high pressures and low temperatures, all gases can be liquefied THE HALOGENS Exist as diatomic molecules – contains two molecules held together by a single covalent bond At room temperature: F2/Cl2= gas Br2= liquid I2= solid Liquid and solid have much higher intermolecular forces and therefore the molecules move less freely THE VAN DER WAALS EQUATION ) x (V – nb) = nRT a and b refer to van der Waals constants (found in a table) MELTING AND BOILING POINTS Melting and boiling points are good indicators of the strengths of intermolecular forces Boiling point = the temperature at which the average kinetic energy of molecular motion balances the attractive energy of intermolecular attractions Normal boiling point is at 1.013 x 10 Pa i.e. boiling point of bromine is 332K – above this, the kinetic energy exceeds the attractive energies created by intermolecular forces and bromine exists as a vapour Vaporisation = liquid  gas Condensation = gas  liquid Freezing point = the kinetic energies decrease and the molecules have too little kinetic energy to slide past one another 5 Normal freezing point is at 1.013 x 10 Pa Stronger intermolecular forces = the higher the freezing point INTERMOLECULAR FORCES Intermolecular forces = tertiary structure Polar groups outside and nonpolar groups inside (when in aqueous environment) DISPERSION FORCES The attractions between negatively charged electron cloud of one molecule and the positively charged nuclei of neighbouring molecules i.e. diatomic halogens A charge imbalance is created, giving Nucleus of one the molecule a molecule attracts slight postive the electron cloud charge at one end of another andat the otherarge Electron cloud is response to this attraction and to minimise electron- electron repulsion Polarisability: the ease of distortion (as distortion creates a temporary polarity)  I2has a higher boiling point and its larger electron cloud means it distorts more readily DIPOLE-DIPOLE FORCES Attractions between the negatively charged end of a polar molecule and the positively charged ends of neighbouring polar molecules  Electrostatic attraction  Uneven distribution of electron density  Polar molecule has a permanent dipole and is arranged asymmetrically  Important over short distances Addition of a carbonyl group (C-O-O) introduces a dipole HYDROGEN BONDING A special case of dipolar forces, requiring a highly polar bond (O-H, N-H or F-H; top right of table) with an electron deficient hydrogen Positive dipole on H + electronegative atom on neighbouring molecule FACTORS EFFECTING INTERMOLECULAR FORCES MOLECULAR STRUCTURE Larger molecule = more polarisable electron cloud = stronger dispersion forces = higher boiling point  More carbon atoms  Tertiary structure of proteins (polar groups outside and nonpolar groups inside) ELECTRONEGATIVITY Relative ability of a bonded atom to attract shared electrons measured on Pauling Scale A polar bond has a difference of > 0.4 in electronegativity SOLUBILITY Maximum amount of solute that will dissolve in a given amount of solvent at a specific temperature (heterogeneous equilibrium) Hydrophobic: will not dissolve Hydrophilic: will dissolve Polar solvent + polar solute Nonpolar solvent + nonpolar solute Solute-solvent interactions > solute-solute interactions for a substance to dissolve SOLIDS INTERMOLECULAR FORCES IN SOLIDS Strong intermolecular forces mean there is no large scale movement of molecules Polar solute dissolves in polar solvent – molecules are attracted to polar water molecules: Lowering of energy (enthalpy) Nonpolar solute dissolves in a nonpolar solvent – dispersion forces are of comparable strength No change of enthalpy but molecules are mixed up (entropy – increased probability) Amorphous – no extended order i.e. glass, some plastics Solids Crystalline – highly ordered i.e. diamond, NaCl Network > metallic > ionic NETWORK SOLIDS Highest melting point  Array of covalent bonds which links each atom to its neighbouring atoms - SiO2(Quartz) - Diamond (covalent bonds between atoms) - Graphite (covalent bonds + dispersion forces between layers) IONIC SOLIDS Cations or anions held together by opposing electrostatic forces  Strong interactions = high melting temperatures  Melting involves separating ions Formulas express relative numbers of cations and ions (infinite structure) Ionic lattice structures: oppositely charged ions pack in different arrangements dictated by their radius and their surface charge density METALLIC SOLIDS Highly electrical and thermal conductivity Highly malleable and ductile Non-directional covalent bonding  Electrons arranged around metal atoms in such a way that they are mobile  Electrons are delocalised  Interactions of metal atoms result in molecular orbitals Alloys: addition of impurities such as carbon or other metals alters properties such as strength and malleability i.e. addition of carbon into iron lattice prevents adjacent atoms from sliding past each other and hardens iron to steel Human history defined in terms of alloys: 1) Stone Age (pre 10,000 BC – 3300 BC) 2) Bronze Age (3300 BC – 500/1000 BC) (copper + tin) 3) Iron (steel) Age (1000 BC – 500 BC/500 AD) MOLECULAR SOLIDS Individual molecules occupy lattice positions (e.g. ice)  Intermolecular forces weaker than covalent forces  Molecules bound by dispersion, dipolar or hydrogen bonds  Generally moderately low melting points Overcoming attractive forces (examples): Sugar dissolves in water = break hydrogen bonds Diamond is crushed = break carbon-carbon covalent bonds Ice melts = break hydrogen bonds Octane boils = overcome dispersion forces Sodium chloride dissolves in water = break ionic bonds Copper wire stretched until it breaks = break covalent interactions in copper lattice ENERGY AND CHEMICAL REACTIONS HEAT AND WORK Heat and work are the only ways that a chemical system can exchange energy with its surroundings 1. Capacity to do work (w) Work done = force x distance 2. The capacity to transfer heat (g) Heat = the process of transfer of thermal energy between two bodies or systems at different temperatures Most chemical reactions absorb or release energy Universe = surroundings + system  Open system: can exchange matter and energy with surroundings  Closed system: can exchange only energy with surroundings  Isolated system: cannot exchange matter or energy with the surroundings  Adiabatic system:
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