Unit 6: Physical Behavior of matter

Definite Composition (Homogeneous) I. Classification of Matter Chart Matter Substance Definite Composition (Homogeneous) Physically Separable Mixture of Substances

I. Classification of Matter Chart Substance Element Contains just a single element Ex: Iron (Fe), Oxygen (O2) Compound Two or more different elements Ex: Water (H2O), Carbon Dioxide (CO2) Chemically Separable

Checks for Understanding A compound differs from an element in that a compound Is homogeneous Has a definite composition Has a definite melting point Can be decomposed by a chemical reaction Which of the following substances cannot be separated by chemical change? Nitrogen (g) Sodium chloride (s) Carbon dioxide (g) Magnesium Sulfate (aq)

Nonuniform; distinct phases I. Classification of Matter Chart Mixture of Substances Homogeneous Uniform throughout (air, salt water, aqueous solutions (aq)) Heterogeneous Nonuniform; distinct phases (soup, concrete)

Check for Understanding A pure substance that is composed only of identical atoms is classified as… A compound An element A heterogeneous mixture A homogeneous mixture 2. A heterogeneous material may be… A mixture Pure substance

II. Particle Diagrams

III. Separating Matter Certain types of matter can be separated using various methods. Monatomic Elements - _____________ be decomposed (broken apart) using _____________ or ______________ means. Diatomic Elements and Compounds (ie – O2 and H2O) – can be decomposed using __________________ only CANNOT CHEMICAL PHYSICAL CHEMICAL MEANS

III. Separating Matter Mixtures – can be separated using ___________________   Filtration – Evaporation – Chromatography – Distillation – PHYSICAL MEANS Separation by particle size Separation by boiling point Separation by polarity Separation by boiling point

Check for Understanding Which of the substances could be decomposed by a chemical change? A) sodium B) aluminum C) magnesium D) ammonia A sample of a material is passed through a filter paper. A white deposit remains on the paper, and a clear liquid passes through. The clear liquid is then evaporated, leaving a white residue. What can you determine about the nature of the sample? What are some of the differences between a mixture of iron and oxygen and compound composed of iron and oxygen? It is a heterogeneous mixture In a mixture the elements are not bonded with each other and can be physically separated. In a compound the elements are bonded and can only be separated through a chemical reaction.

Think about this What would a particle diagram look like for each of these phases? What happens to the spacing and speed of particles at each of the phases? SOLID LIQUID GAS

IV. Forms of Mechanical Energy Kinetic Energy Energy of movement (similar to temperature) (how fast atoms are moving) Potential Energy Stored energy (energy of position) More spread out (gas) = High PE Closer together (solid) = Low PE

V. Heating and Cooling Curves (animation) ENDOTHERMIC ABSORBED Heating Curve: ___________ - Energy is being ________  gas l  g  liquid s  l s  g  solid Sublimation (video)- Solid changes directly to a gas Heating Curve Animation

V. Heating and Cooling Curves AB BC CD DE EF Kinetic Energy Potential Energy Phase Con-stant Con-stant ↑ ↑ ↑ Con-stant Con-stant Con-stant ↑ ↑ gas solid l  g boiling s  l melting liquid

Check for Understanding A substance begins to a melt. What happens to the potential and kinetic energy? PE increase, KE stays the same 2. The temperature of a substance refers to what type of energy? Kinetic energy 3. How does the speed and space of water molecules compare when in a liquid phase to a gas phase Molecules move faster and more spread out in gas phase

V. Heating and Cooling Curves EXOTHERMIC RELEASED Cooling Curve: ___________ - Energy is being ________  gas g  s g  l  liquid l  s  solid Deposition - Gas changes directly to a solid

V. Heating and Cooling Curves AB BC CD DE EF Kinetic Energy Potential Energy Phase Con-stant Con-stant ↓ ↓ ↓ Con-stant Con-stant Con-stant ↓ ↓ g  l condensing solid gas liquid l  s Freezing

Check for Understanding As a substance condenses, what happens to its potential and kinetic energy? PE decreases, KE stays the same 2. What phase is a substance in when it has its highest kinetic energy? gas 3. How does the speed and space of water molecules compare when in a liquid phase to a solid phase Molecules move slower and are closer together in solid phase

Vi. Temperature vs. Heat Amount of energy transferred from one substance to another Average kinetic energy of its particles (how fast they’re moving) Joules (J) or Calories (cal) 1 cal = 4.18 J Celsius (oC) or Kelvin (K) (K = oC + 273) T q

Vi. Temperature vs. Heat K = oC + 273 K = Kelvin oC = degrees Celsius Temperature Scales (See Ref. Tabs.): K = oC + 273 K = Kelvin oC = degrees Celsius   Convert: 200 degrees Celsius to Kelvin Law of Conservation of Energy: Heat Transfer: K = oC + 273 K = 200oC + 273 = 473 K Energy (heat) cannot be created or destroyed. Energy (heat) can be TRANSFERRED. HEAT ALWAYS MOVES FROM WARMER OBJECTS TO COLDER OBJECTS

Check for Understanding You wake up in the morning and your barefoot touches the ceramic floor and it feels cold. Explain which way heat is being transferred. Heat moves from your body (warm) to the floor (cold) You are cooking pasta in a boiling metal pot of water. You grab the metal handles with your bare hands (ouch!). Explain which way heat is being transferred. Why do you feel cold after you get out of a hot shower. (link) Heat moves from metal handles (warm) to your hands (cold).

vii. Endothermic and exothermic (revisited) Energy is either absorbed or released in chemical reactions Remember:   - Breaking bonds is ____________________ - Heat is ___________ the reaction from the surroundings - Ex) heat + Br2  Br + Br - Creating bonds is _________________ - Heat is ___________ the reaction into the surroundings - Ex) N+ N N2+ energy ENDOTHERMIC ENTERING EXOTHERMIC EXITING

vii. Endothermic and exothermic (revisited) Where does the heat come from (or go to)? ______________________   For exothermic reactions, heat (energy) leaves the reaction and moves __________________________. Therefore, making the surrounding temperature _________________ Exothermic Chemical Equation: The surroundings into the surroundings warmer Reactant(s)  Product(s) + HEAT

vii. Endothermic and exothermic (revisited) For endothermic reactions, heat (energy) leaves the surroundings and moves __________________________. Therefore, making the surrounding temperature _________________ Endothermic Chemical Equation: Into the reaction colder Reactant(s) + HEAT  Product(s) ENDOTHERMIC REACTION VIDEO

vii. Endothermic and exothermic (revisited) Enthalpy: It is a measure of the heat released or absorbed in a reaction. If heat is released then the reaction is an exothermic reaction, heat will be on the products side and the ΔH will be negative. If heat is absorbed then the reaction is an endothermic reaction, heat will appear on the reactants side and the ΔH will be positive. (see Table ____ on Reference Tables). Examples: Tell whether each of the following reactions are endothermic or exothermic (you may have to use table I). Then tell whether the temperature of the surroundings increases or decreases as a result.   Endo/Exothermic _______Surroundings __________ C6H12O6 + 6O2  6CO2 + 6H2O + heat _______________ I Exo (-∆H) warmer

vii. Endothermic and exothermic (revisited)   Endo/Exothermic _______Surroundings ____________C6H12O6 (s) + 6O2(g)  6CO2 (g) + 6H2O + heat ____________ 2. __________ 2H2O + 484 kJ  2H2 + O2 _______________ 6. __________ 2KClO3(s)  2KCl(s) + O2(g) + heat _______________ 7. __________ H+(aq) + C2H3O2(aq) + heat  HC2H3O2(l) _______________ exo(-∆H) warmer endo(+∆H) colder exo(-∆H) warmer endo(+∆H) colder

VIiI. Vapor Pressure ______ VP = _________________________ Example:   ______ VP = _________________________ Example: Factors for Vapor Pressure 1. 2. Pressure a liquid “feels” pushing it to evaporate (turn to gas) HIGH Evaporates easily Ethanol has higher VP than H2O, so it evaporates more easily Strength of intermolecular force: molecules held together by dipole-dipole (polar) have lower VP than Van der Waals (nonpolar) (H2O does not evaporate as fast as methane (CH4)) Temperature/Pressure: increase temp., increase VP

VIiI. Vapor Pressure Vapor Pressure of Four Liquids (See Table _____)   Questions: 1. What is the vapor pressure of propanone at 35 oC? _______ 2. What temperature does water boil at if the pressure is 70 kPa? ______ 3. What is the normal boiling point of ethanoic acid? __________ Sublimation: Example: “Dry Ice” CO2 (s)  CO2 (g) H Measured in: kPa (101.3 kPa = 1 atm = 760 mmHg = 760 torr.) 48 kPa 90 degrees C 117 degrees C Occurs because solids have very weak IMF (usually Van der Waals). They go directly from s  g and have HIGH VP

Check for Understanding How could you change the boiling point temperature of a substance without adding anything to the substance? You spill a glass of water on the floor. How could you get the water to evaporate faster, without using a mop or something to soak it up?   Change your altitude (higher altitude (lower pressure), lower boiling point) Increase the temperature of the room. Spread out the puddle (increase surface area).