Phase Changes

10 Key excerpts on "Phase Changes"

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  Hands-On Science and Technology for Ontario, Grade 5

  An Inquiry Approach

  • Jennifer E. Lawson, Jennifer Lawson(Authors)
  • 2020(Publication Date)
  • Portage & Main Press

    (Publisher)

  Changes of state are physical changes, because no new materials are created. Consider the behaviour of water. As solid ice, water molecules have little energy. They can only vibrate enough to stay in a regular arrangement. Add heat energy, and they begin to vibrate more. Eventually, the molecules can overcome their intermolecular attractions enough to slide over one another. The water has become liquid. With more heat added, the temperature of the water increases, and the molecules become more active. If enough heat is added, some of the water molecules can escape the pull of their neighbours and become gaseous. The water has begun to boil. With continued addition of heat, all the water molecules will convert to gas. Temperature Molecular Motions Molecular Interactions Solid Liquid Gas The changes described are physical changes, as the characteristics and properties of the water molecules have not changed. Each molecule is still two hydrogen atoms bonded to one oxygen atom. It is possible to capture the gaseous water and condense it back into its liquid state, and freeze the liquid into a solid. 21 st Century Competencies Critical Thinking and Communication: Students will examine water as it changes state, and understand what is occurring as the phase change occurs. Materials ■ ice cubes ■ electric frying pan ■ aluminum baking pan ■ oven mitts ■ clear, plastic cup ■ student journals ■ access to a local body of water ■ access to a freezer ■ Activity Sheet: Changing States of Water (3.6.1) ■ Learning-Centre Task Card: How Important Is Precipitation to Life on Earth? (3.6.2) ■ resources about rain, snow, and the water cycle ■ markers ■ paper (ledger size or larger) ■ sticky notes ■ concept web (from lesson 1) ■ Science and Technology Glossary (3.1.2) Activate Begin the activity by reviewing the three states of matter.

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  Visualizing Everyday Chemistry

  • Douglas P. Heller, Carl H. Snyder(Authors)
  • 2015(Publication Date)
  • Wiley

    (Publisher)

  When you heat a solid, though, its chemical particles absorb energy. When the temperature of the solid reaches intermolecular forces Attractive forces that exist between molecules in close proximity. Distinguishing states of matter • Figure 6.1 All common objects are either a solid, a liquid, or a gas. Some substances, like water, can be in any state, depending on factors such as temperature. Ask Yourself Which of the three common states of matter—solids, liquids, or gases—(a) maintain their own shapes, no matter what container holds them? (b) maintain their own volumes, no matter what container holds them? a. Solid - - retains own shape b. Liquid - - adopts shape of container c. Gas - adopts volume and shape of container fixed volume fixed volume 165 move about freely within the bulk of the material, we observe that the solid melts to a liquid (Figure 6.2). Melting and other changes in state are examples of a physical change. its melting point, the movements of its particles become sufficiently vigorous to tear them away from their neighbors and out of their fixed positions. As they begin to What happens when solids melt • Figure 6.2 When we heat a solid to its melting point, individual particles, such as molecules, gain enough energy to break out of their fixed positions. This is what happens when ice melts to form water. WHAT A CHEMIST SEES States of Matter at the Molecular Level Water is found in its three physical states in this image of geothermal activity in Yellowstone National Park. In each state of matter, water molecules differ in their arrangements and movements. melting point The temperature at which a solid is transformed into a liquid. physical change A transformation of matter that occurs without any change in chemical composition. Gas In water vapor, water molecules move about at high speeds and at relatively large distances from one another.

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  The Basics of Physics

  • Richard L. Myers(Author)
  • 2005(Publication Date)
  • Greenwood

    (Publisher)

  Focusing on the three common states of matter, six different Phase Changes are pos- sible. These are summarized in Table 6.2. The most familiar phase change occurs between the solid and liquid states (freezing/ melting) and the liquid and gaseous states (vaporization/condensation). Figure 6.3 illustrates the heating curve for water and depicts how temperature changes as heat is added to a substance. Initially, when ice is below its freezing point, heating will increase the ice's temperature just as with other solids. Once the ice reaches its melt- ing point of 0° C (at 1 atm pressure), the ice starts to melt. At this stage both solid and liquid exist. The heating curve plateaus at the melting point even though heat is still being added to the ice-liquid mixture. It may seem strange that the temperature does not increase even though energy is added, but Heat 89 Table 6.2 Phase Changes Melting Solid-^Liquid Sublimation Solid^Gas Vaporization Liquid^Gas Freezing Liquid—>Solid Deposition Gas^Solid Condensation Gas^Liquid this can be explained by the fact that the added energy is used to convert ice to liquid water. The energy supplied at the melting point is not being used to increase the ran- dom kinetic energy of the water molecules, but to overcome intermolecular attractions between water molecules. The energy sup- plied at the melting point goes into breaking the water molecules free from the crystal- line structure, resulting in the less-structured liquid state. As long as ice is present, any energy added goes into causing the phase change. The heat necessary to melt the ice is termed the heat of fusion for water, and its value is 6.0 kJ per mole of water. The heat of fusion of a substance is a measure of how much energy is required to convert a solid into a liquid. It would be expected that solids that are tightly held together would have high heats of fusion.

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  Thermal Physics

  Entropy and Free Energies

  • Joon Chang Lee(Author)
  • 2011(Publication Date)
  • WSPC

    (Publisher)

  Chapter 12 Phases and Phase Transitions 12.1 Introduction When the temperature is high, the constituent particles act independently and freely, and therefore they can be treated as a collection of ideal gas molecules, ideal spins or ideal rods, etc. Alas, the gas molecules actually interact with one another as do the magnetic spins, and the interactions can no longer be ignored when the temperature is not sufficiently high. Those molecules come close to each other and turn into a different phase, namely liquid. We will not study interacting molecules per se , but we will find a way to take a glimpse into the thermal world of interacting molecules and observe how the free energy is altered by the molecular interactions. We will then observe how the altered free energy alters molecular behavior. Here, molecules are effectively cornered so that the only way to minimize the free energy is to change their basic structure. Yes, they decide that they are better off by coming closer to each other and to condense into a liquid. They are just like school children running around on their play ground. When the bell rings, they too know that they are better off by stop running and come to their classrooms. What a turnaround! That is also the trend for magnetic spins. They too interact with each other, and when the temperature is low, they act together and orient in the same direction. Such an orderly behavior at low temperatures, as opposed to the random and disorderly behavior at high temperatures, is our theme from now on. The same maximum entropy or the minimum free energy principle apply, but under the changed condition, the orderly behavior is the best way to maximize their entropy or to minimize their free energy. 237 238 Phases and Phase Transitions 12.2 Liquid State and Solid State When we studied the gas phase with the ideal gas model, the spatial part of the multiplicity had nothing but the incredibly simple V N factor.

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  Phase Transitions in Materials

  • Brent Fultz(Author)
  • 2014(Publication Date)
  • Cambridge University Press

    (Publisher)

  Throughout much of this text, the detailed interactions between the entities of matter are replaced with simplifying assumptions that facilitate mathemat- ical modeling. Sometimes the essence of the phase transition is captured well with such a simple model. Other times the discrepancies prove interesting in themselves. Perhaps surprisingly, the same mathematical model reappears in explanations of phase transitions involving very different types of matter. A phase transition is an “emergent phenomenon,” meaning that it displays features that emerge from interactions between numerous indi- vidual entities, and these large-scale features can occur in systems with very different microscopic interactions. The study of phase transitions has become a respected field of science in its own right, and Chapter 20, for example, presents some concepts from this field that need not be grounded in materials phenomena. 1.2 Atoms and materials An interaction between atoms is a precondition for a phase transition in a material (and, in fact, for having a material in the first place). Atoms interact in interesting ways when they are brought together. In condensed matter there are liquids of varying density, and numerous types of crystal structures. Magnetic moments form structures of their own, and the electron density can show spatial modulations. In general, chemical bonds are formed when atoms are brought together. The energy of interatomic interactions is dominated by the energy of the electrons, which are usually assumed to adapt continuously (“adiabat- ically”) to the positions of the nuclei. The nuclei, in turn, tend to position themselves to allow the lowest energy of the material, which means that nuclei move around to let the electrons find low-energy states. Nevertheless, once we know the electronic structure of a material, it is often possible to understand many properties of a material, especially its chemical, electronic, magnetic, and optical properties.

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  Chemistry for Today

  General, Organic, and Biochemistry

  • Spencer Seager, Michael Slabaugh, Maren Hansen, , Spencer Seager, Spencer Seager, Michael Slabaugh, Maren Hansen(Authors)
  • 2021(Publication Date)
  • Cengage Learning EMEA

    (Publisher)

  Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. 186 Chapter 6 In some instances, solids cannot be changed into liquids, or liquids into gases, by heating. The atoms making up the molecules of some solids acquire enough kinetic energy on heating to cause bonds within the molecules to break before the solid (or liquid) can change into another state. This breaking of bonds within molecules changes the composition of the original substance. When this decomposition occurs, the original substance is said to have decomposed. This is why cotton and paper, when heated, char rather than melt. 6.15 Energy and the States of Matter Learning Objective 13 Do calculations based on energy changes that accompany heating, cooling, or changing the state of a substance. A pure substance in the gaseous state contains more energy than in the liquid state, which in turn contains more energy than in the solid state. Before we look at this, note the following relationships. Kinetic energy, the energy of particle motion, is related to heat. In fact, temperature is a measurement of the average kinetic energy of the particles in a system. Potential energy, in contrast, is related to particle separation distances rather than motion. Thus, we conclude that an increase in temperature on adding heat corresponds to an increase in kinetic energy of the particles, whereas no increase in temperature on adding heat corresponds to an increase in the potential en- ergy of the particles. Now let’s look at a system composed of 1 g of ice at an initial temperature of 220°C. Heat is added at a constant rate until the ice is converted into 1 g of steam at 120°C. The atmospheric pressure is assumed to be 760 torr throughout the experiment. The changes in the system take place in several steps, as shown in Figure 6.21.

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  Basic Physical Chemistry

  The Route to Understanding

  • E Brian Smith(Author)
  • 2012(Publication Date)
  • ICP

    (Publisher)

  9 The States of Matter 9.1 Gases, liquids and solids Matter can exist in many forms but, most commonly, we identify three distinct states: gas, liquid and solid. The state with the lowest Gibbs free energy is the stable form of matter at any particular temperature. At low temperatures, the solid with the most negative energy is the most stable form. At high temperatures, the gaseous state with the maximum randomness prevails. At intermediate temperatures, the liquid state has the lowest free energy. If the Gibbs free energy is plotted against the temperature at constant pressure, since d G = V d P − S d T , the slopes of the lines are (∂G/∂T) P = − S . The gaseous phase with the highest entropy has the largest negative slope and will have the lowest free energy at high temperatures (Fig. 9.1). The solid phase with the lowest slope has the lowest free energy at low temperatures. The transition from one phase to another occurs where the lines intersect and where the free energies are equal. Then, G = 0 and H = TS , giving, at the melting point of the solid, fus S = fus H T fus and, at the boiling point of the liquid, vap S = vap H T vap . We can represent the equilibrium of the phases as a function of temperature and pressure as shown in Fig. 9.2 for the phase equilibria in water. Such diagrams are called phase diagrams . The lines represent the pressures and temperatures at which two phases are in equilibrium. The line AB represents the vapour pressure of liquid water and AC gives the vapour pressure of ice. AD is the melting curve on which the solid and liquid are at equilibrium. The point where all three curves meet, A, is called the triple point , which, for water, is at 273.16 K and 10.61 kPa (4.58 mm Hg) pressure, the only pressure and temperature at which the three phases can co-exist in equilibrium. 199 200 | Basic Physical Chemistry G T T fus T vap Solid Liquid Gas Fig.

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  Engineering Materials Science

  Properties, Uses, Degradation, Remediation

  • H McArthur, D Spalding(Authors)
  • 2004(Publication Date)
  • Woodhead Publishing

    (Publisher)

  2 STATES OF MATTER AND PHYSICAL CONSTANTS 2.1 INTRODUCTION There are three states of matter, a gas or vapour, a liquid (the most common one being water) and a solid (which may be amorphous or crystalline). These three states (or phases) are shown by water as the solid ice, the liquid water and as the vapour steam. In this Chapter we focus on water, as this compound is responsible for a wide range of building defects and deterioration mechanisms (Chapters Sand 6). 2.2GASES Gas molecules exlubit constant translocational movement (visualised as billiard balls colliding) and fill the available volume. Gas molecules are modelled as having principally translocational kinetic energy and a relatively small amount ofvibrational kinetic energy (movement which stretches the interatomic bonds, e.g, Cl-Cl, 0-0, etc.). An increase in temperature increases the translocational kinetic energy to a greater extent than the vibrational kinetic energy. A decrease in temperature results in the molecules of a gas coming into closer proximity, so that bonding can occur and the gas (water vapour) condenses to form a liquid. 2.2.1 Ideal Gas Laws If one mole of a gas is confined in volume V at temperature T, the pressure exerted by the gas, p, obeys the relationship pV =RT, where R is the universal gas constant(= 8414 J/kmol.K) and Tis absolute temperature (K). For n moles of the gas, pV = nRT, where n = m!M and mis the mass of the gas of molecular weight M. This equation is the equation of state of a perfect gas, also known as the ideal ( perfect) gas law. 2.2.2 Fundamental Kinetic Theory Equation for a Gas The kinetic energy possessed by a body by virtue of its motion is called the kinetic energy of the body. Suppose a gas molecule of mass m moving with velocity u is bought to rest in a distance s by a constant retarding force F (Figure 2.1 a). The original kinetic energy of the gas molecule is equal to Fs, and this must therefore be the work done in bringing the molecule to rest.

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  200 Science Investigations for Young Students

  Practical Activities for Science 5 - 11

  • Martin Wenham(Author)
  • 2000(Publication Date)
  • SAGE Publications Ltd

    (Publisher)

  Section 6.5 ).

  Most simple materials, i.e. those which are not mixtures, exist either as solids, liquids or gases with no intermediate or transitional states between them, though solids may soften as they are heated. Under the same conditions the same material will always melt and freeze at the same temperatures (its melting- and freezing-points) and the change in either direction will be sudden, though it may not always appear to be. Children may associate ‘freezing’ only with changes which take place at temperatures they experience as ‘cold’, but scientifically freezing is the reverse of melting, regardless of the temperature at which it takes place.

  Activity 6.2.1

  Observing melting and freezing

  The most useful material for introductory investigations into melting and freezing is candle (paraffin) wax. Ordinary white (i.e. uncoloured) candles should be used, since materials used to colour candles or crayons can interfere with observation of melting and freezing. Any sample of paraffin wax will melt and freeze at a definite temperature, usually between 50°C and 60°C and the change between the two states can easily be seen.

  These and other investigations which involve the heating and cooling of small amounts of material (see also Activity 6.2.3 ) can effectively be carried out using small food trays or containers made of fairly thick aluminium foil, floating on water.

  Observing melting and freezing of wax

  Equipment and materials: White (i.e. uncoloured) candle; knife; small aluminium foil food containers; plastic bowl; electric kettle or other means of heating water; cold water; thermometer, range − 10 to 110°C; a long-handled spoon.

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  Chemical Thermodynamics at a Glance

  • H. Donald Brooke Jenkins(Author)
  • 2008(Publication Date)
  • Wiley-Blackwell

    (Publisher)

  The point B corresponds to the boiling temperature, T b at the (constant) pressure P involved. From the above rationale (and without having made any prior or additional assumptions) we see that our analysis leads – via thermodynamic argument – to a prediction that pure materials existing in their normal phases exhibit melting and boiling phenomena, as is observed experimentally. 64 Variation of G with T for Solid, Liquid and Gaseous Phase G s < G l < G g at T = 0 K G S T T A solid solid liquid Pressure = P liquid P ressure = P gas gas 0 0 C E F T m T b T m T b T B ( ∂ G s / ∂ T ) P = − S l ( ∂ G l / ∂ T ) P = − S l ( ∂ G g / ∂ T ) P = − S g G s G l G g This area corresponds to figure 19.3 discontinuity at phase transition s l Δ fus S = Δ fus H T m T b Since S s < S l the solid G s versus T curve is flatter than the G l versus T curve resulting in an intersection at T m Since S g >> S l or S S the gas curve dips down the most and intersects with the liquid curve at T b large discontinuity at phase transition l g Δ vap S = Δ vap H Figure 21.1 Graph of G g , G l and G s and of S g , S l and S s versus temperature, T , at constant pressure, P (d P = 0) for a normal substance existing in three simple phases (solid, liquid and gas (vapour)). As a caveat to the considerations involving the solid phase it should be mentioned that some solids undergo mesomorphic changes be-fore melting (although these have been assumed not to occur in the solids so far discussed). 21.2 Variation of Entropy, S with Temperature, T for the Solid, Liquid and Gaseous Phases of a Pure Substance, at Constant Pressure (d P = 0) In Figure 21.1, placed directly below the graph of G g , G l and G s versus T , we sketch a graph of the corresponding entropy func-tions: S g , S l and S s versus temperature, T whilst mapping from the upper to the lower graph the phase stabilities as predicted at the various values of temperature, T , in the upper graph.

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