Name: gas pressure patterns in kinetic molecular theory
Uploaded: 2020-09-03T13:38:27.000Z
Duration: 2 min 34 s
Channel: JoVE

K12 Standards

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Chemistry

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PS.Chm.7.3

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Kinetic Molecular Theory and Gas Laws Explain Properties of Gas Molecules

ملاحظة خصائص الغاز المختلفة،كما عبرت عنه قوانين الغاز المختلفة المستمدة من بويل،تشارلز،و جاي-لوساك وأفوجادرو،تتبع من الناحية المفاهيمية النظرية الجزيئية الحركية. ينتج الضغط الذي يمارسه الغاز من تأثير الجسيمات المتحركة باستمرار على جدران الإناء الخاصة به. تقليل حجم الإناء،مع الحفاظ على عدد المولات ودرجة الحرارة ثابتة،تقرب جزيئات الغاز من بعضها البعض،مما يقلل التباعد بين الجسيمات.في هذا الحجم الأصغر،كثافة الغاز وتواتر الاصطدام تواتر اصطدام الجزيءبالجدار-تزداد. لذلك،يزداد الضغط الذي يمارسه الغاز. العلاقة عكسية بين الضغط والحجم بموجب قانون بويل.إضافة المزيد من مولات الغاز إلى الإناء عند درجة حرارة ثابتة تزيد من كثافة الغاز،وبالتالي،تردد الاصطدام. للحفاظ على الضغط الأولي،يجب أن يتوسع الحجم. هذه العلاقة المباشرة بين الحجم والمولات يتم تقديمها بموجب قانون أفوجادرو.الآن،لنضع في اعتبارنا عدد المولات تبقي ثابتة ودرجة الحرارة ترتفع. لأن متوسط الطاقة الحركية لجزيئات الغاز تتناسب مع درجة الحرارة،تتصادم الجسيمات بشكل متكرر وقوي. إذا كان الحجم ثابتًا في هذه الأثناء بينما تزداد درجة حرارة،فإنّكثافة الغاز و تردد الاصطدام سيزدادا،وبالتالي سيزداد الضغط أيضًا.العلاقة المباشرة بين الضغط ودرجة الحرارة يتم تقديمها بموجب قانون جاي-لوساك. من ناحية أخرى،إذا كان الضغط يجب أن يبقى ثابتًا إلى جانب عدد ثابت من المولات،فإنّاي ارتفاع في درجة الحرارة يجب ان يصاحبه زيادة في الحجم لتوزيع الإصطدامات على مساحة أكبر. هذه العلاقة المباشرة بين الحجم ودرجة الحرارة معطاة بواسطة قانون تشارلز.أخيرًا،وفقًا للنظرية الجزيئية الحركية،جزيئات الغاز لا تجتذب أو تتنافر مع بعضها البعض. في خليط من الغازات المختلفة،تعمل المكونات بشكل مستقل و تظل ضغوطهم الفردية لا تتأثر بوجود غاز آخر. الضغط الكلي للخليط هو،بالتالي،مجموع الضغوط الفردية الجزئية.هذا هو قانون دالتون.

النظرية الجزيئية الحركية وقوانين الغاز

إن اختبار النظرية الجزيئية الحركية (KMT) والمسلمات الخاصة بها هو قدرتها على شرح ووصف سلوك الغاز. من الممكن أن تستمد قوانين الغاز المختلفة (قوانين بويل’، وتشارلز’، ولاس جاي لوساك’، وأفاوجادرو’، ودالتون’) من افتراضات قانون KMT، الذي دفع الكيميائيين إلى الاعتقاد بأن الافتراضات التي تستند إليها هذه النظرية تمثل بدقة خصائص جزيئات الغاز. <h4>تشرح النظرية الجزيئية الحركية (KMT) سلوك الغازات </h4>مع الإشارة إلى أن ضغط الغاز يتم بواسطة جزيئات الغاز سريعة الحركة ويعتمد بشكل مباشر على عدد الجزيئات التي تصل إلى منطقة وحدة من الجدار لكل وحدة من الوقت، يشرح الكيومينتانغ مفاهيميا سلوك الغاز كما يلي: <ul><li>قانون جاي لوساك’s: إذا ارتفعت درجة الحرارة فإن متوسط السرعة والطاقة الحركية لجزيئات الغاز تزداد. إذا كان الحجم ثابتاً، فإن زيادة سرعة جزيئات الغاز تؤدي إلى حدوث تصادمات أكثر تكراراً وقوة مع جدران الحاوية، وبالتالي زيادة الضغط. ويطلق على هذا أيضاً أمونتونز&rsquo قانون.</li><li>قانون تشارلز&rsquo: إذا ارتفعت درجة حرارة الغاز، فإنه بالإمكان الحفاظ على  ضغط ثابت فقط عندما يزداد الحجم الذي يشغله الغاز. سيؤدي ذلك إلى مسافات متوسطة أكبر تقطعها الجزيئات للوصول إلى جدران الوعاء الحاوي، بالإضافة إلى زيادة مساحة سطح الجدار. وسوف تقلل هذه الظروف من تردد تصادمات الجزيء والجدار وعدد التصادمات لكل منطقة وحدة، والتي تعمل آثارها مجتمعة على موازنة تأثير قوى التصادم المتزايدة بسبب الطاقة الحركية الأكبر في درجة الحرارة الأعلى. </li><li>قانون بويل’: إذا انخفض حجم الغاز، فإن منطقة جدار الوعاء الحاوي تنخفض، ويزداد تردد التصادم بين الجزيء والجدار، وكل من الأمرين يزيد من الضغط الذي يمارسه الغاز. </li><li>قانون أفوجادرو’: في ظل ضغوط ودرجة حرارة ثابتة، فإن تكرار وقوة التصادمات بين الجزيء والجدار الثابت. وفي ظل هذه الظروف، تتطلب زيادة عدد الجزيئات الغازية زيادة نسبية في حجم الوعاء الحاوي من أجل تقليل عدد التصادمات في منطقة الوحدة الواحدة للتعويض عن زيادة تكرار التصادمات. </li><li>قانون دالتون’: بسبب المسافات الكبيرة بينهما، فإن جزيئات غاز واحد في خليط ما تقوم بقصف جدران الوعاء الحاوي بنفس التردد، سواء كانت غازات أخرى موجودة أم لا، والضغط الإجمالي لمزيج من الغاز يساوي مجموع الضغوط (الجزئية) للغازات الفردية. </li></ul>هذا النص مقتبس من <a href="https://openstax.org/books/chemistry-2e/pages/9-5-the-kinetic-molecular-theory">Openstax, Chemistry 2e, Section 9.5: The Kinetic-Molecular Theory.</a>

اختبار النظرية الجزيئية الحركية (KMT) وافتراضاتها هو قدرتها على شرح ووصف سلوك الغاز. قوانين الغاز المختلفة ...

Education

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الكيمياء

واحد

Gases

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KEY TERMS AND CONCEPTS

المصطلحات والمفاهيم الأساسية

gas pressure patterns in kinetic molecular theory

Kinetic molecular theory explains gas pressure by linking it to the motion of gas molecules. Gas pressure comes from rapidly moving molecules hitting the walls of a container. It depends on how many molecules strike a unit area in a given time.

The theory also helps explain the main gas laws. In Gay-Lussac’s law, also called Amontons’s law, raising temperature increases the average speed and kinetic energy of gas molecules. If volume stays constant, the molecules collide with the walls more often and more forcefully, so pressure increases.

Charles’s law shows what happens when temperature rises at constant pressure. The gas must expand to keep pressure steady. That larger volume means molecules travel farther before reaching the walls, and the wall area increases too. These changes lower collision frequency and the number of collisions per unit area, which balances the stronger collisions caused by the higher kinetic energy.

Boyle’s law describes the effect of smaller volume. When gas volume decreases, the container wall area also decreases. Molecule-wall collisions become more frequent, and the pressure rises. Avogadro’s law says that at constant pressure and temperature, the collision rate and collision force stay constant. If the number of gas molecules increases, the volume must increase in proportion so the collisions per unit area can drop enough to offset the added molecules.

Dalton’s law applies to gas mixtures. Because gas molecules are far apart, one gas in a mixture strikes the container walls with the same frequency whether other gases are present or not. The total pressure of the mixture equals the sum of the partial pressures of the individual gases. These gas laws show how the assumptions of kinetic molecular theory fit the behavior of gases.

Gas Pressure Patterns in Kinetic Molecular Theory

gas-pressure-patterns-in-kinetic-molecular-theory

definition and measurement of pressure: atmospheric pressure, barometer, and manometer

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Pressure Units and Tools for Gas Measurement

Pressure is the force applied to a given area, and gas pressure comes from gas molecules colliding with surfaces. Each collision is tiny, but many collisions over a small area can create high pressure. Pressure increases when force increases or when the area gets smaller. It decreases when force drops or when the area gets larger. Pressure is measured in pascals, or Pa, in the SI system. One pascal equals 1 newton per square meter, and a newton is a unit of force equal to 1 kg·m/s2. A pascal...

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How Gas Pressure, Volume, and Temperature Change

Gas pressure, volume, temperature, and amount are linked by several gas laws. Scientists found these relationships by studying how a confined gas changes when one variable is held constant and another changes. These patterns apply to an ideal gas, which is a useful model for gases under low pressure and high temperature. Gay-Lussac’s law, also called Amontons’s law, connects pressure and temperature. If gas is trapped in a rigid container, the volume and number of moles stay constant. When the...

applications of the ideal gas law: molar mass, density, and volume

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Ideal Gas Law at STP and Gas Density

The ideal gas law links pressure, volume, number of moles, and temperature. It helps explain how gas volume changes when pressure or temperature changes. At standard temperature and pressure, or STP, the values are 273.15 K and 1 atm. At STP, one mole of an ideal gas occupies about 22.4 L. This is called the standard molar volume. The gas can be hydrogen, oxygen, argon, or carbon dioxide, and the volume is still about 22.4 liters. Half a mole occupies 11.2 L, and 2 moles occupy 44.8 L. The...

mixtures of gases: dalton's law of partial pressures and mole fractions

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Gas Pressure in Mixtures: Partial Pressures

Gas pressure in mixtures is described by partial pressure and mole fraction. When gases do not chemically react, each gas in the mixture acts as if it were alone in the container. The pressure from one gas is called its partial pressure. For a mixture of three gases, A, B, and C, the total pressure equals the sum of their partial pressures: PA, PB, and PC. This is Dalton's law of partial pressures. It states that the total pressure of a mixture of ideal gases is the sum of the pressures from...

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Stoichiometry of Gases with the Ideal Gas Law

Chemical stoichiometry links the amounts of reactants and products in a chemical reaction. For gases, those amounts can also be measured by volume. If a gas’s volume, pressure, and temperature are known, the ideal gas equation can be used to find the number of moles present. The same idea works in reverse. If the moles of a gas are known, its volume can be found at any temperature and pressure. The transcript gives an example using 8.88 g of gallium reacting with excess hydrochloric acid.

basic postulates of kinetic molecular theory: particle size, energy, and collision

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Kinetic Molecular Theory and Gas Behavior

Kinetic molecular theory explains gas behavior at the particle level. It gives a microscopic model for what happens when pressure or temperature changes. Gas laws describe how macroscopic properties change, but they do not explain why those changes happen. The theory was put into a complete form by Rudolf Clausius in 1857. It was built from hundreds of experimental observations about gases. Those observations helped explain the different gas laws through a few basic postulates. One key idea...

kinetic molecular theory and gas laws explain properties of gas molecules

Kinetic molecular theory explains gas pressure by linking it to the motion of gas molecules. Gas pressure comes from rapidly moving molecules hitting the walls of a container. It depends on how many molecules strike a unit area in a given time. The theory also helps explain the main gas laws. In Gay-Lussac’s law, also called Amontons’s law, raising temperature increases the average speed and kinetic energy of gas molecules. If volume stays constant, the molecules collide with the walls more...

kinetic molecular theory: molecular velocities, temperature, and kinetic energy

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Gas Particle Speed Distributions and Temperature

Gas particle speed distributions and temperature are key ideas in the kinetic molecular theory. This theory explains the gas laws by describing how gas molecules move, collide, and share energy. Gas molecules do not all move at the same speed. A gas sample has an average speed and an average kinetic energy, but individual molecules move at different speeds. Their collisions are elastic, which means momentum is conserved. Because so many molecules collide, the overall speed distribution stays...

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real gases: effects of intermolecular forces and molecular volume deriving van der waals equation

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Real Gases: When the Ideal Gas Law Breaks Down

Real gases are gases that do not always follow the ideal gas law, PV = nRT, especially when pressure is high or temperature is low. The ideal gas model assumes gas particles have no attractive or repulsive forces and take up negligible space compared with the container. At room temperature and 1 atm or less, many gases behave close to ideally. When conditions change, the measured relationships between pressure, volume, and temperature can drift away from the ideal prediction. One way to see...

key terms and concepts

Explore how kinetic molecular theory explains gas pressure, temperature, volume changes, and the gas laws of Boyle, Charles, Gay-Lussac, Avogadro, and Dalton.

Gas Pressure Patterns in Kinetic Molecular Theory | JoVE K12

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Vapor Pressure

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Clausius-Clapeyron Equation

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Heating and Cooling Curves

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Energy Basics

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Chemical changes and their accompanying ...

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Internal Energy

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Quantifying Heat

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Quantifying Work

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Enthalpy

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PS.Chm.7.2

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Kinetic Molecular Theory: Molecular Velocities, Temperature, and Kinetic Energy

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Behavior of Gas Molecules: Molecular Diffusion, Mean Free Path, and Effusion

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Real Gases: Effects of Intermolecular Forces and Molecular Volume Deriving Van der Waals Equation

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Gas Pressure Patterns in Kinetic Molecular Theory

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