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chem test
| Question | Answer |
|---|---|
| what are the 4 gas laws on the first page that we supposedly already new | Gases take the volume and shape of the container gases are highly compresible gases will mix evenly and completely when confined to the same container (Homogeneous Gases have low densities when compared with liquids and solids |
| How big are gas particles? | Gas particles are super tiny compared to the space between them. 👉 They’re so small we can pretend they’re just little dots. |
| How do gas particles move? | They are always moving randomly. 👉 They move in straight lines until they hit something. 👉 When they crash, they bounce off and don’t lose energy. |
| Do gas particles attract or repel each other? | No. 👉 They don’t pull toward each other. 👉 They don’t push away from each other. |
| What does temperature measure in a gas? | Temperature measures how fast the particles are moving. 👉 Higher temperature = particles move faster. 👉 Lower temperature = particles move slower. ⚠️ Always use Kelvin for gas laws. |
| What causes gas pressure? | Pressure comes from gas particles hitting the walls of the container. 👉 More hits = higher pressure 👉 Harder hits = higher pressure |
| Does an ideal gas actually exist? | No. 👉 Ideal gases are just a model (a pretend idea) to help us predict behavior. |
| What gases behave most ideally? | Small, light gases like: 👉 Hydrogen (H₂) 👉 Helium (He) Small particles = act more “perfect.” |
| What happens when gas particles collide? | 👉 They hit each other and the container walls. 👉 They bounce off. 👉 Collisions change their speed and direction. |
| Why don’t real gases behave ideally? | Because real gases: 👉 Take up space (they have volume) 👉 Attract each other a little Ideal gas assumes none of that. |
| When do real gases act MOST like ideal gases? | 👉 Low Pressure 👉 High Temperature Memory trick: More Ideal = Low P, High T |
| Why does low pressure make gases more ideal? | Low pressure → big volume Particles are far apart They can’t feel attraction |
| Why does high temperature make gases more ideal? | High temperature → high kinetic energy Particles move fast They overpower any attraction |
| What are the 4 main gas variables? | P = Pressure V = Volume T = Temperature n = Moles |
| What is Pressure? | Force hitting an area. In gases: particles hitting container walls. 1 atm = 760 mmHg = 760 torr = 101.3 kPa |
| What is Volume? | How much space something takes up. |
| What is Temperature in gas laws? | Average kinetic energy (how fast particles move). ⚠️ MUST use Kelvin Formula: K = °C + 273 |
| What is a mole (n)? | Measures the number of particles in a gas. |
| According to the kinetic theory, the most significant difference between gases and liquids is | the distance between the particles |
| An ideal gas is an imaginary gas that | conforms to all of the assumptions of the kinetic molecular theory |
| One difference between a real gas and an ideal gas is that in a real gas | the particles exert attractive forces on each other. |
| 4) The conditions under which real gases most resemble ideal gases are | low P and high T |
| A gas at low temperature does not behave like an ideal gas because | the KE of the particles is too low |
| The case in which the kinetic theory does not hold true for gases can be explained by | forces between molecules |
| What does Dalton’s Law state? | The total pressure of a mixture of gases equals the sum of the partial pressures of each individual gas. |
| Dalton’s Law formula | P Total= P1 + P2 + P3 + ... |
| Dalton’s Law Example Concept | If total pressure = 278 kPa O₂ = 112 kPa H₂ = 101 kPa Nitrogen pressure = 278 − 112 − 101 = 65 kPa |
| What does Avogadro’s Principle say? | Equal volumes of gases at the same temperature and pressure contain the same number of particles (same number of moles). |
| What is molar volume at STP? | At STP, 1 mole of gas = 22.4 L |
| What is effusion? | Gas molecules escaping through tiny holes in a container. |
| What is diffusion? | Gas spreading out (like perfume in a room). |
| Central concept of Graham’s Law | Large particles move slower than small particles when escaping. |
| If a gas diffuses 1/50 as fast as hydrogen (2.02 g/mol), what does that mean? | Slower rate = larger molar mass. Use Graham’s Law and solve for unknown molar mass. |
| How to find density of CO at given T and P | d = MP / RT Plug in molar mass of CO, pressure, temperature, and R. |
| How to find molar mass at STP if density is known | M = dRT / P At STP → P = 1 atm, T = 273 K |
| Density formula | d=m/v |
| Ideal Gas Law formula | PV=nRT |
| How do you find the partial pressure of a gas using moles? | Add total moles Divide gas moles by total moles Multiply by total pressure |
| when it says sm like theres 2 flasks with the same amount of volume whats the same in the________ | number of molecules |
| What happens to temperature if volume doubles and pressure stays constant? | Use Charles’s Law. If volume doubles → temperature doubles (in Kelvin). Example: 200 K → 400 K Answer: 200 K to 400 K |
| How do you know if two gases have the same number of particles? | Use the Ideal Gas Law relationship. If volume is the same, then: from the table: Ne → 1.00 / 300 CH₄ → 1.00 / 300 |
| What is the relationship between pressure and Kelvin temperature at constant volume? | Pressure and temperature increase together in a straight line. |
| What is the relationship between pressure and volume at constant temperature? | PV = constant |
| What happens when Kelvin temperature increases at constant pressure? | Molecules move faster → kinetic energy increases Volume increases |
| If temperature increases and pressure stays constant, how does volume change? | Volume increases slightly |
| Which gas diffuses fastest? | Smaller molar mass → faster diffusion |
| What property of gas molecules is directly related to temperature? | Average Kinetic Energy |
| According to the Kinetic Theory of Gases, what happens to molecular motion when temperature increases? | Molecules move faster (average kinetic energy increases). |
| At constant volume, what is the relationship between the speed of gas molecules and the pressure they exert? | Faster molecules = more frequent and forceful collisions with walls = higher pressure. |
| Which graph best represents the relationship between pressure and average kinetic energy at constant volume? | A direct linear relationship (as KE goes up, Pressure goes up). Think: Line going upwards from left to right. |
| Key Concept: Kinetic Theory of Gases. | Pressure is caused by molecules colliding with container walls. Average Kinetic Energy is related to temperature. At constant volume, higher KE means higher Pressure |
| What is Boyle’s Law (verbal)? | At constant temperature, pressure and volume are inversely proportional. |
| What is Charles’s Law (verbal)? | At constant pressure, volume is directly proportional to temperature. |
| What is Gay-Lussac’s Law (verbal)? | At constant volume, pressure is directly proportional to temperature. |
| What conditions cause gas laws to break down? | Gas laws stop working well at: • High pressure • Low temperature Particles get too close together and start attracting each other. |
| What is Graham’s Law? | The rate of diffusion or effusion of a gas is inversely proportional to the square root of its molar mass. Smaller molar mass → faster gas |
| What does a Boyle’s Law graph look like? | Pressure vs Volume is an inverse curved relationship (hyperbola). |
| What does a Charles’s Law graph look like? | Volume vs Temperature is a straight line increasing when temperature is in Kelvin. |
| What does a Gay-Lussac graph look like? | Pressure vs Temperature is a straight line increasing when temperature is in Kelvin. |
Created by:
Fernando.Quezada