States of matter
51 flashcards covering States of matter for the ACT Science section.
States of matter refer to the distinct forms that substances can take based on how their particles are arranged and move. The main states are solid, liquid, and gas, with solids having tightly packed particles that vibrate in place, liquids featuring particles that slide past one another, and gases consisting of particles that spread out and move freely. These states can change through processes like melting or evaporation, influenced by factors such as temperature and pressure, and they play a key role in understanding physical properties and chemical behaviors in the world around us.
On the ACT Science section, states of matter typically appear in questions involving data interpretation, graphs of phase changes, or basic properties like density and boiling points. Common traps include mistaking one state's characteristics for another or overlooking how variables like heat affect transitions, so watch for misleading answer choices. Focus on recognizing patterns in experiments and applying concepts to real-world scenarios to tackle these questions effectively.
A helpful tip: Always check how temperature and pressure impact state changes in graphs.
Terms (51)
- 01
Solid
A solid is a state of matter where particles are tightly packed in a fixed arrangement, giving it a definite shape and volume, and particles vibrate in place.
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Liquid
A liquid is a state of matter that has a definite volume but takes the shape of its container, with particles close together but able to move past one another.
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Gas
A gas is a state of matter with no definite shape or volume, where particles are far apart and move freely, filling the available space.
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Plasma
Plasma is a state of matter consisting of ionized gas with free-moving electrons and ions, commonly found in stars and fluorescent lights, and it conducts electricity.
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Kinetic theory of matter
The kinetic theory of matter states that all matter is made of tiny particles in constant motion, and the state of matter depends on the average kinetic energy and distance between particles.
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Intermolecular forces
Intermolecular forces are the attractions between molecules that determine the state of matter, being strongest in solids and weakest in gases, affecting properties like boiling and melting points.
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Melting
Melting is the phase change from solid to liquid when a substance absorbs enough heat to overcome intermolecular forces, occurring at the melting point.
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Freezing
Freezing is the phase change from liquid to solid as a substance loses heat and intermolecular forces pull particles into a fixed structure.
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Evaporation
Evaporation is the process where liquid molecules escape into the gas phase at the surface, even below the boiling point, due to gaining sufficient kinetic energy.
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Condensation
Condensation is the phase change from gas to liquid when gas molecules lose kinetic energy and come together, forming droplets on a cool surface.
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Sublimation
Sublimation is the phase change directly from solid to gas without passing through the liquid state, as seen in dry ice turning into carbon dioxide gas.
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Deposition
Deposition is the phase change directly from gas to solid, bypassing the liquid state, such as water vapor forming frost on a window.
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Boiling point
The boiling point is the temperature at which a liquid turns into a gas throughout the liquid, occurring when vapor pressure equals atmospheric pressure.
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Melting point
The melting point is the specific temperature at which a solid becomes a liquid, depending on the substance's intermolecular forces and pressure.
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Vapor pressure
Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid or solid phase, increasing with temperature and affecting boiling.
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Phase change
A phase change is a transformation of matter from one state to another, such as solid to liquid, driven by changes in temperature or pressure.
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Heat of fusion
Heat of fusion is the amount of energy required to change a substance from solid to liquid at its melting point without changing temperature.
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Heat of vaporization
Heat of vaporization is the energy needed to convert a liquid to gas at its boiling point, accounting for the breaking of intermolecular forces.
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Boyle's Law
Boyle's Law states that for a given amount of gas at constant temperature, pressure and volume are inversely proportional, meaning if volume decreases, pressure increases.
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Charles's Law
Charles's Law states that for a fixed amount of gas at constant pressure, volume is directly proportional to temperature in Kelvin, so heating a gas expands it.
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Gay-Lussac's Law
Gay-Lussac's Law states that for a given amount of gas at constant volume, pressure is directly proportional to temperature in Kelvin.
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Combined Gas Law
The Combined Gas Law relates pressure, volume, and temperature of a gas, stating that P1V1/T1 = P2V2/T2 for a fixed amount of gas.
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Ideal Gas Law
The Ideal Gas Law, PV = nRT, relates pressure, volume, number of moles, gas constant, and temperature for an ideal gas, assuming no molecular volume or interactions.
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Avogadro's Law
Avogadro's Law states that equal volumes of different gases at the same temperature and pressure contain an equal number of molecules.
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Dalton's Law of Partial Pressures
Dalton's Law states that the total pressure of a mixture of non-reacting gases is the sum of the partial pressures of the individual gases.
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Graham's Law of Effusion
Graham's Law states that the rate of effusion of a gas is inversely proportional to the square root of its molar mass, explaining why lighter gases escape faster.
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Pressure
Pressure is the force exerted per unit area by gas particles colliding with the walls of their container, measured in units like atm or pascals.
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Volume
Volume is the space occupied by a gas, which can change with pressure and temperature, and is a key variable in gas laws.
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Absolute temperature
Absolute temperature, measured in Kelvin, is the thermodynamic temperature scale starting at absolute zero, where molecular motion theoretically stops.
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Standard temperature and pressure (STP)
STP is defined as 0 degrees Celsius and 1 atm pressure, where one mole of an ideal gas occupies 22.4 liters.
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Molar volume
Molar volume is the volume occupied by one mole of a substance, which for an ideal gas at STP is 22.4 liters.
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Diffusion
Diffusion is the process by which gas particles spread out from an area of higher concentration to lower concentration due to their random motion.
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Effusion
Effusion is the escape of gas molecules through a tiny hole into a vacuum, with the rate depending on the gas's molar mass as per Graham's Law.
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Real gases vs. ideal gases
Real gases deviate from ideal gas behavior at high pressures and low temperatures due to molecular volume and intermolecular forces, unlike ideal gases.
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Phase diagram
A phase diagram is a graph showing the conditions of temperature and pressure at which a substance exists as solid, liquid, or gas, including phase boundaries.
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Triple point
The triple point is the unique temperature and pressure on a phase diagram where a substance can exist in solid, liquid, and gas phases in equilibrium.
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Critical point
The critical point is the temperature and pressure beyond which a gas cannot be liquefied, marking the end of the liquid-gas boundary on a phase diagram.
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Supercritical fluid
A supercritical fluid is a state of matter above its critical point, with properties between those of a liquid and a gas, used in applications like extraction.
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Crystalline solids
Crystalline solids have a regular, repeating arrangement of particles in a lattice structure, leading to definite melting points and shapes.
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Amorphous solids
Amorphous solids lack a regular internal structure, behaving more like liquids on a molecular level, with no sharp melting point, like glass.
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Viscosity
Viscosity is a measure of a liquid's resistance to flow, depending on intermolecular forces and temperature, with higher viscosity in thicker liquids.
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Surface tension
Surface tension is the property of a liquid that allows it to resist an external force due to cohesive forces between molecules at the surface.
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Density
Density is the mass per unit volume of a substance, which varies between states of matter, being highest in solids and lowest in gases.
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Compressibility
Compressibility is the measure of how much a substance's volume decreases under pressure, with gases being highly compressible unlike solids and liquids.
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Brownian motion
Brownian motion is the random movement of particles in a fluid due to collisions with fast-moving molecules, evidence of particle motion in liquids and gases.
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Effect of temperature on states
Increasing temperature generally adds kinetic energy to particles, causing matter to change from solid to liquid to gas, overcoming intermolecular forces.
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Effect of pressure on states
Increasing pressure can force a gas into a liquid or solid state, as seen in the compression of gases or the boiling point of liquids.
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Strategy for gas law problems
To solve gas law problems, identify the given variables, choose the appropriate law, convert temperatures to Kelvin if needed, and solve for the unknown using the formula.
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Common trap: Temperature units
A common error is using Celsius instead of Kelvin in gas laws, which invalidates results since the laws require absolute temperature.
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Example of Boyle's Law
In Boyle's Law, if a gas at 2 atm and 4 L is compressed to 2 L, the pressure becomes 4 atm, showing the inverse relationship.
A balloon shrinks when squeezed, increasing internal pressure.
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Example of Charles's Law
In Charles's Law, if a gas at 273 K and 1 L is heated to 546 K at constant pressure, the volume doubles to 2 L.
A hot air balloon rises as the air inside expands with heat.