States of Matter and Energy Dr. Walker

States of Matter Phase – A part of matter that is uniform Solid (lowest energy) Liquid Gas Plasma (highest energy) Found in stars, lightning

State Changes Enthalpy = Energy

Intermolecular Forces Intramolecular Forces – bonding between atoms to achieve octet Ionic Bonding Covalent Bonding

Intermolecular Forces Forces between separate molecules As energy is added (higher temperature), molecules have more motion (higher KE) and overcome intermolecular forces to change state These forces differ with the type of compound (polar vs. nonpolar) Determines physical properties (bp/mp, density, etc.)

Intermolecular Forces Dipole-Dipole Forces Interactions between polar molecules Results from unequal sharing About 1% of strength of ionic bonds “partial charges” – shows unequal sharing between the two atoms

Review… Polar molecules Molecule with difference in electronegativity between atoms Not symmetric Unequal sharing of electrons (shown with d symbol

Hydrogen Bonding Special dipole-dipole attraction Hydrogen covalently bonded to highly electronegative elements (N, O, F) has a higher than normal d+ charge Bond strength is higher than other dipole-dipole attractions VERY IMPORTANT in biology Links strands of DNA Involved in protein folding (3D shape)

T A C G Thymine hydrogen bonds to Adenine Cytosine hydrogen bonds to Guanine C G

Hydrogen Bonding Gives water unique properties Solid is less dense than liquid (why ice floats)

London Dispersion Forces (aka Van der Waals Forces) Movement of electrons in an atom create temporary dipole Attraction between nonpolar molecules Much weaker than dipole interactions

London Dispersion Forces https://i.ytimg.com/vi/P2CF2JBZAY8/maxresdefault.jpg

Intermolecular Forces vs. Melting/Boiling Points Ionic Networks – Strongest interaction Covalent Bonding Hydrogen Bonding (dipole forces) – Water, Compounds with N, O London Dispersion Forces – hydrocarbons (Weakest Interaction)

Bonding Strength vs. MP Stronger interaction = higher boiling/melting point Order of increasing melting points: London Dispersion (hydrocarbons) < Hydrogen bonding and dipole interactions (covalent compounds with N, O, halogens) < Ionic bonding compounds

Comparing Intermolecular Forces vs. Boiling Point Material Symbol/Formula Boiling Point (oC) Nonpolar Covalent Hydrogen H2 -253 Methane CH4 -164 Polar Covalent Ammonia NH3 -33 Water H2O 100 Ionic Sodium Chloride NaCl 1413 Magnesium Oxide MgO 2826

Vapor Pressure Pressure created by gas molecules above the surface of a liquid Volatile liquids (alcohol, gas, nail polish remover) have a high vapor pressure Nonvolatile liquids (water) have a low vapor pressure Solutions have a lower vapor pressure (remember colligative properites! More vapor = Higher vapor pressure!

Vapor Pressure Inversely proportional to intermolecular (IM) forces Weak intermolecular forces allows more gas to form, resulting in higher vapor pressure Polar compounds – higher IM forces – less gas formation – higher bp Nonpolar compounds – low IM forces – more gas formation – lower bp

Vapor Pressure When vapor pressure = atmospheric pressure, a substance will boil Lower vapor pressure – further from 1 atm – more heat needed to increase pressure Higher vapor pressure – closer to 1 atm – less heat needed to increase pressure

Inc. polarity Inc. IM forces Dec. VP Inc. MP Notice – as the boiling points increase… 1) ..the vapor pressures start off lower 2) …the polarity increases (propane, ethanol, water, ethanoic acid) 3) …the IM forces increase (with polarity)

Vocabulary Enthalpy Entropy Total Energy of a system Disorder of a system

Measuring Energy What do we measure energy in? The Joule is the SI (metric) system unit for measuring heat: The calorie is the heat required to raise the temperature of 1 gram of water by 1 Celsius degree 1 BTU is the heat required to raise the temperature of 1 pound of water by 1 F

Energy and Heating/Cooling What is the relationship between heat and temperature? Heat – transferred energy Temperature – measure of kinetic energy If you transfer the same amount of heat in different substances, the temperature change will not be the same The temperature change is governed by…

Specific Heat Capacity Determines temperature changes vs. heat transfer The energy required to raise 1 g of material by 1 degree celsius E or E = mCDT change in E = Energy m = mass C = specific heat DT = temperature Temperature scale doesn’t matter, 1 C = 1 K Different substances have different specific heat capacities Metals tend to be low (>1 J/g x K) Water is very high (4.18 J/g x K) Reason why a blacktop gets hotter much faster than a pool on a sunny day!

Example A 4.0 g sample of glass was heated from 274 K to 314 K and absorbed 32 J of heat. What is the specific heat?

Example A 4.0 g sample of glass was heated from 274 K to 314 K and absorbed 32 J of heat. What is the specific heat? Specific Heat = 32 J/(4.0 g x [314-274]) 32 J/(4.0 g x 40 K) 0.20 J/g x K

Example Water has a specific heat of 4.18 J/g C. How much energy is required to raise the temperature of a 1000 g sample by 5 C?

Example Water has a specific heat of 4.18 J/g C. How much energy is required to raise the temperature of a 1000 g sample by 5 C? E = mCDT E = (1000 g)(4.18 J/g C)(5 C) = 20900 J or 20.9 kJ

Energy of Phase Changes When compounds change phase, the enthalpy (energy) and entropy (disorder) changes Plasma, Gas have highest entropy Solid has lowest entropy Disorder – think of freedom of molecular movement Gases have highest entropy Solids have lowest entropy

What About During A Phase Change? When a substance is at the melting or boiling point, the temperature does not change until the phase change is complete For example, water boils at 100 oC. Temperature can’t rise to 101 oC until the entire sample in steam! The energy added is breaking intermolecular forces, not to moving the molecules.

Energy of Phase Changes Notice that during a phase change, energy is added, but the temperature does not change

Molar Heat of Fusion The energy required to melt one mole of a solid at its melting point Molar Heat of Fusion for ice = 6.00 kJ/mole Molar Heat of Fusion for methane (CH4) = 1592 J/mole These are per mole, not per gram. You could convert from grams if necessary…

Molar Heat of Fusion How much energy is required to change 6 g of ice to water at 0 C (melting point)?

Molar Heat of Fusion How much energy is required to change 6 g of ice to water at 0 C (melting point)? You have to find the moles of water first by dividing by molar mass 6 g/(18 g/mole) = 0.33 moles 0.33 moles x 6.00 kJ/mole = 1.98 kJ of energy

Molar Heat of Vaporization Energy required to convert 1 mole of liquid to 1 mole of gas at its boiling point Water = 40.65 kJ/mole Methane (CH4) = 8.19 kJ/mole Ethanol (C2H6O) = 38.6 kJ/mole More polar compounds have a higher heat of vaporization (more energy required to break intermolecular forces)

Molar Heat of Vaporization How much energy is required to change 90 g of ethanol (C2H6O) from liquid to gas at its boiling point?

Molar Heat of Vaporization How much energy is required to change 90 g of ethanol (C2H6O) from liquid to gas at its boiling point? 90 g/(46 g/mole) = 1.96 moles 1.96 moles x 38.6 kJ/mole = 75.7 kJ

Energy of a Heating Curve

Relationships Between Physical Properties http://www.dlt.ncssm.edu/tiger/diagrams/phase/EnergyChange_and_Phases.gif

Phase Diagrams Definition Graph of pressure vs. temperature Can tell the phase of a substance at a given point Every compound has a different phase diagram

Phase Diagram of Water Point A – Triple Point All three phases at equilibrium at this temperature and pressure

Phase Diagram of Water Point C – Critical Point Above this temperature and pressure, the substance CANNOT exist as a liquid

Phase Diagram of Water Atmospheric Pressure = 1.0 atm The temperatures on the phase lines are the melting point and boiling point

Interpret A Phase Diagram

Energy and Reactions All Chemical Reactions either release or absorb energy – it just isn’t always noticeable Explosions – chemical reaction with a large amount of energy released

Energy and Reactions Exothermic – Reaction process that gives off energy Reactants  Products + Energy Combustion reactions – give off a great deal of energy in the form of fire Neutralization reactions – acid/base reactions typically get fairly hot

Energy and Reactions Endothermic – Reaction process that absorbs energy Reactants + Energy  Products

Reaction Coordinates Exothermic Reaction Reaction Progress starts with reactants Ends with products Energy (H) at beginning higher than Energy at end Negative Enthalpy (-DH) Add acid + base as a demo, feel the beaker

Reaction Coordinates Endothermic Reaction Energy (H) at end higher than Energy at beginning Positive Enthalpy (+ DH) http://cimg2.ck12.org/datastreams/f-d%3Ad07bc599a7310d9fd6fbfea4b44861cb93fdcd869142ecfb61bc2b4b%2BIMAGE%2BIMAGE.1

Quick Example What kind of reaction is this? Is the DH positive or negative?

Activation Energy Activation Energy – The minimum amount of energy required to start a reaction

Catalysts A catalyst lowers the activation energy necessary to start a chemical reaction Changes the rate of a chemical reaction without being consumed Point 2 demonstrates a catalyst on a reaction coordinate

Quick Example Is this reaction endothermic or exothermic? What is the DH of this reaction?

Common Catalysts Catalytic Converter Enzymes Made of precious metals (Pt, Rh, Pd) 2 CO + O2 2 CO2 Enzymes Proteins which speed up metabolic processes Cells could not function without enzymes Enzyme malfunctions result in disease Lactase (lactose intolerance)

What Can Affect Reaction Rates? Concentration – More molecular collisions, faster rate Temperature – Higher temperature, faster reactions Molecules move faster, more molecular collisions

What Can Affect Reaction Rates? Surface Area – More exposed area (think sugar cube vs. granulated sugar), faster reactions Pressure (gas) – More pressure = more molecular collisions, faster reactions