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Matter – States of Matter, Properties and Changes

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  1. Matter – States of Matter, Properties and Changes

  2. Matter • Matter takes up space and has mass • Matter is made of atoms, usually chemically bonded into molecules • Exists in different states

  3. States of Matter • There are 4 states of matter: solid, liquid, gas, and plasma • State of a sample of matter depends on the kinetic energy of the molecules or atoms in the sample • Kinetic energy is the energy of moving things

  4. States of Matter • There are 4 states of matter: solid, liquid, gas, and plasma • State of a sample of matter depends on the kinetic energy of the molecules or atoms in the sample • Kinetic energy is the energy of moving things

  5. Solids • Solids have a definite shape and a definite volume • The atoms and molecules that make a solid, vibrate in place but do not move around

  6. Kinetic Theory of Matter, Solids • Particles in solid matter are held close together by forces between them • Particles vibrate but don’t have enough energy to move out of position

  7. Liquids • Liquids have a fixed volume, but take the shape of the container in which they are found • The atoms and molecules that make a liquid can flow around each other

  8. Kinetic Theory of Matter, Liquids • Particles in liquid matter are held close together by forces between them • Particles are close enough so that liquid matter has a definite volume • Particles have enough energy to move over and around each other

  9. Gases • Gases have neither a definite shape nor volume • They take the shape of their container

  10. Kinetic Theory of Matter, Gases • Particles of a gas have enough energy to separate completely from one another • Particles of a gas are not close together so they can be squeezed into a smaller space • Particles have enough energy to move in all directions until they have spread evenly throughout their container

  11. Plasma • Plasma is a gaslike mixture of positively and negatively charged particles • They have so much energy that they collide violently and break apart into charged particles • Found in lightning bolts, neon signs, Northern lights, and stars

  12. It is made of electrons and positive ions that have been knocked apart by collisions at very high temperatures or in situations where the matter has absorbed energy • Least common state of matter on Earth but is the most common state of matter in the universe, because stars are made of matter in the plasma state

  13. Thermal Expansion • Almost all matter expands when it gets hotter and contracts when it cools • When matter is heated the particles move faster, vibrate against each other with more force • Particles spread apart slightly in all directions and the matter expands

  14. This effect happens in solids, liquids, and gases • Examples are the liquid in a thermometer and expansion joints in roads and buildings

  15. Changes of State • When matter gains or loses energy, it can change from one state to another • Different states of matter correspond to different amounts of energy, these amounts are specific to particular kinds of matter • Temperature can be used to measure the amounts of energy present in the matter

  16. Change of state terms • Boiling: liquid changes to a gas • Freezing: liquid changes to a solid • Condensing: gas changes to a liquid • Melting: solid changes to a liquid • Evaporating: liquid changes to a gas (but a temperatures lower than the boiling point) • Subliming: solid changes into a gas without becoming liquid (opposite of sublimation is deposition)

  17. Change of state temperatures • Boiling point: temperature at which a liquid becomes a gas, this temperature is an identifiable characteristic for different substances • Melting point: temperature at which a solid becomes a liquid

  18. Substances condense or boil at their boiling point, depending on whether energy is being added or taken away • Substances melt or freeze at their melting point, depending on whether energy is being added or taken away

  19. Phase changes • Transitions between solid, liquid, and gaseous phases typically involve large amounts of energy compared to the energy needed to change the temperature of a solid or liquid or gas. • It takes lots of energy to change states (temperature stays even until states are completely changed).

  20. If heat were added at a constant rate to a mass of ice to take it through its phase changes from solid to liquid water and then to steam, the energies required to accomplish the phase changes would lead to plateaus in the temperature vs time graph.

  21. Boiling point elevation • http://group.chem.iastate.edu/Greenbowe/sections/projectfolder/flashfiles/propOfSoln/colligative.html interactive boiling point and freezing point changes • Adding solute to water increases its boiling point, the solute interacts with the water and energy must be added to overcome the interactions so that the water can then change from a liquid to a gas

  22. Freezing point depression • http://group.chem.iastate.edu/Greenbowe/sections/projectfolder/flashfiles/propOfSoln/colligative.html interactive boiling point and freezing point changes • Adding solute to water decreases its freezing point, the solute interacts with the water and energy must be removed to overcome the interactions so that the water can then change from a liquid to a solid

  23. Thermal Energy • Thermal energy is the total energy of the particles in a material • Thermal energy includes the kinetic energy of the particles (their motion or vibration) • Thermal energy also includes the potential energy of the particles (energy due to forces acting within or between the particles)

  24. Heat • Heat is the name given to thermal energy that moves or is transferred • In many things that you read, heat and thermal energy are used interchangeably

  25. Movement of Heat • Heat moves from areas of greater heat (more thermal energy) to areas of lesser heat (less thermal energy)

  26. Temperature • Temperature is the measure of the average kinetic energy of the particles that make up a sample of matter. • As the particles move faster, the temperature rises • As the particles slow down, the temperature falls

  27. Temperature scales - Fahrenheit • First scale developed • Water melts/freezes at 32°Fand boils at 212°F • Salt water melts/freezes at 0°F, body temperature was 96°F and degrees were divided into 12s and then into 8s between these two

  28. Temperature scales - Celsius • Celsius scale based on 100 degrees between freezing and melting of water • Water melts/freezes at 0°Cand boils at 100°C

  29. Temperature scales - Kelvin • Important scale used in most of science • Based on a single point (absolute zero) which is given a value of 0 degrees. • From there, the scale increases by degrees that are the same size as Celsius degrees.

  30. Temperature scales - Kelvin • It is a scale that is based on energy content, rather than on arbitrary temperature values like the other two scales (based on water). • Water freezes at the value 273.15 K and boils at 373.15 Kelvin.

  31. Absolute Zero • 0 on the Kelvin scale • Point at which all particle motion stops • Matter has no thermal energy at absolute zero

  32. Law of Conservation of Energy • Law of Conservation of Energy – Energy is neither created nor destroyed. It can change forms.

  33. Specific Heat Capacity • Physical property of matter • Relates to a substance’s ability to absorb heat • Also called specific heat

  34. Specific Heat Capacity • Specific heat capacity of a substance is the amount of energy (Joules) required to raise the temperature of 1 gram of the substance by 1 °C Specific heat capacity =

  35. Objects with low specific heat capacities heat up more quickly than objects with high specific heat capacities. • It takes less energy to raise their temperatures • They also transfer their heat more quickly so they cool down faster

  36. Water has a fairly high specific heat capacity, 4.184 J/g °C • This means it takes a lot of energy to raise the temperature of water 1 °C compared to the amount of energy it takes to heat something with a lower specific heat capacity • Example: Iron (0.45 J/g °C) need much less energy to change its temperature

  37. Heat conductors and insulators • Objects with low specific heat capacities are better conductors of heat • Objects with high specific heat capacities are better insulators because they don’t heat up as quickly

  38. Calorimeter • An insulated container that prevents a chemical reaction from gaining heat from its surroundings or losing heat to its surroundings

  39. Using a calorimeter to calculate specific heat capacity • Calorimeter experiments to calculate specific heat capacities of objects use the Law of Conservation of Energy and the known specific heat capacity of water • When a heated object is placed in a cup of cold water, the heat will move from the object to the water • When the temperature stops changing, the temperature of the object and water are now the same

  40. Using a calorimeter to calculate specific heat capacity • When the temperature stops changing, the temperature of the object and water are now the same • Energy transferred to the water is equal to the energy transferred from the object

  41. Calculations: Known: Specific heat capacity of water = 4.184 J/g °C Energy transferred to water = mass of water (g) x Temp change (°C) x 4.184 J/g °C Specific heat capacity of object =