Physical Properties of Matter

Matter

• What is it?

• What is it not?

Anything that has mass and takes up space.

Seeing something is not enough to classify it as matter, energy/waves are not matter.

Matter can contain, emit, and absorb energy but energy cannot contain matter.

Body

* What is it?

A bunch of matter together.

Particle

• What is it?

• Examples

A single unit, could be compounds or singular elements. This is the safe word, it is never wrong.

Examples: NO2, O2, O.

Molecule

• What is it?

• Examples and non-examples

A single unit of a covalent compound only.

Examples: H2O and O2.

CaCO3 is __not__ an example.

Atom

• What is it?

• Examples

One unit of a single element. Specific to the number of elements, not the type of element.

Examples: N is one, H2O is three, CaCO3 is five.

Ion

* What is it?

A charged particle.

Physical Property

• What is it?

• Examples

Describes observable features of a substance.

Examples: State of matter, colour, shine, malleability, rigidity, melting point, and boiling point.

Chemical Property

• What is it?

• Examples

Describes how the substance will interact with something else.

Examples: Reactivity and solubility.

Physical Change

• What is it?

• Examples

Reversible, no new product is formed.

Examples: Melting, boiling, bending.

Chemical Change

• What is it?

• Exampes

Not reversible, always forms a new product.

Examples: Lighting on fire (combustion) and cooking.

How can you tell the difference between a change and a property?

A property is something that the substance can do (i.e. gallium __can__ melt in your hand) whereas a change is something the substance did do/is doing (i.e. gallium __is__ melting in your hand).

Pure Substance

• What type of matter?

• What is it made of?

Matter that __cannot__ be physically separated. It is made up of one type of element or compound (constant composition).

Element

• What type of matter?

• What does this mean?

A pure substance that __cannot__ be chemically decomposed. A single unique atom.

Compound

• What type of matter?

• What does this mean?

A pure substance that __can__ be chemically decomposed. Made up of two or more atoms.

Mixture

• What type of matter?

• How is it formed?

Matter that __can__ be physically separated, an impure substance. Formed when two or more elements/compounds are present without being chemically bonded or reacted.

Homogenous Mixture

• What type of matter?

• What is another name?

• What is it made of?

A mixture that __is__ uniform throughout. Also called a solution. It is made up of ionic compounds when naturally occurring.

Heterogenous Mixture

• What type of matter?

• What is it made of?

A mixture that __is not__ uniform throughout. It is made up of covalent compounds when naturally occurring.

Colloid

• What type of matter?

• What does this mean and why?

A heterogenous mixture with __very small__ particles. Very small particles means less gravity acting on them, so the mixture settles slower.

Suspension

• What type of matter?

• What does this mean and why?

A heterogenous mixture with anything other than very small particles. Larger particles means more gravity acting on them, so the mixture settles quicker.

Closed Container

* What is it?

A theoretical container that is air tight and lets no light, energy, air, or matter in or out.

Alloy

• What is it?

• What are some physical properties?

• Example

Something composed of many metals. They have really high IMFs and really high boiling points.

Example: brass.

Why can only ionic substances become solutions/homogenous mixtures?

Water pulls apart ionic compounds (NaCl becomes Na+ and Cl-) but covalent compounds are in “bubbles” (C6H12O6 remains C6H12O6). Covalent compounds have no ions to split so it stays together.

What is the difference between dissolving and dissociating?

Dissolving is when something is broken into its particles and dissociation is when an ionic compound breaks into its ions.

Examples: NaCl is put into water and breaks into Na+ and Cl-, it is both dissolving and dissociating. C6H12O6 is put into water and remains as C6H12O6, it is only dissolving.

What does (s), (l), (g), and (aq) mean in a chemical reaction?

(s) means solid state.

(l) means liquid state.

(g) means gaseous state.

(aq) means aqueous, which is dissolved in water.

Intermolecular Forces (IMFs)

• What are they?

• Are they a physical/chemical property/change and why?

The forces of attraction that exist between all atoms within a substance. They are physical properties because they determine the melting and boiling points.

When do IMFs change and when do they stay the same?

* Examples

All pure substances have unique IMFs meaning they do not change for that substance, even if it changes phase. IMFs are different between different pure substances.

Example: H2O’s IMFs do not change as it switches from a solid to a liquid to a gas, but H2O’s IMFs are different than NaCl’s IMFs.

How are IMFs related to phase changes?

• Example

• What is the significance of gas particle energy and its IMFs?

Since IMFs are constant, the energy (any type) of a particle is what changes its state.

Example: Solid particles have low energy so their IMFs have more control over the particles, so they stay closer together. Liquid particles have more energy so their IMFs are stretched and the particles move farther away from each other. Gas particles have high energy so their IMFs are overcome and the particles have free motion and move farthest away from each other.

Note: Gas particles’ IMFs “disappear” because their energy is so much greater than them, but the IMF’s influence returns when the particle energy decreases again and they change back into a liquid.

Volume

* What is it?

How much 3D space an object occupies.

Shape

* What is it?

The physical boundaries of an object.

Particle Arrangement

* What is it?

The space between particles.

Particle Energy

* What is it?

The amount of radiation that a particle consumes.

Solid

• Volume

• Shape

• Particle Arrangement

• Particle Energy vs. IMFs

• Naturally Occurring?

**Volume:** Set

**Shape:** Set

**Particle Arrangement:** In a set position close together, vibrate in place, gentle collisions with one another.

**Particle Energy vs. IMFs:** IMFs are greater than particle energy.

**Naturally Occurring?:** Yes

Liquid

• Volume

• Shape

• Particle Arrangement

• Particle Energy vs. IMFs

• Naturally Occurring?

**Volume:** Set

**Shape:** Takes the shape of the container.

**Particle Arrangement:** Some ability to move, collisions have more energy which creates more space between particles.

**Particle Energy vs. IMFs:** IMFs are about equal to particle energy.

**Naturally Occurring?:** Yes

Gas

• Volume

• Shape

• Particle Arrangement

• Particle Energy vs. IMFs

• Naturally Occurring?

**Volume:** Will expand as much as possible.

**Shape:** Takes the shape of the container.

**Particle Arrangement:** Free motion, collisions are high energy which generates lots of space between particles.

**Particle Energy vs. IMFs:** IMFs are less than particle energy.

**Naturally Occurring?:** Yes

Plasma

• Volume and Shape: Conditions

• Particle Arrangement: Difference Between Fully and Partially Ionized

• Particle Energy vs. IMFs

• Naturally Occurring?: Conditions

**Volume and Shape:** __IF__ we had a container that could sustain/withstand 1,000,000°C these would be the same as a gas (expands as much as possible and takes the shape of the container).

**Particle Arrangement:** Fully ionized (does not exist on Earth) is gas elevated to a state where all particles are ionized (electrons are in an excited state, valence electrons jump up and hold energy). The sun is a fully ionized plasma. Partially ionized (exists on Earth) is when some/few particles are ionized, but most remain in a gaseous state. This can occur when a gas is electrocuted.

**Particle Energy vs. IMFs:** IMFs are less than particle energy.

**Naturally Occurring?:** Instantaneously (lightning), partially in TVs and neon signs (“partial plasma”), but not sustainable on Earth.

Energy

• What is it?

• What forms does it come in?

• What binds energy?

The potential to do work. Can be in the form of heat, light, motion (kinetic), etc.

Mass is needed to bind energy.

Why are plasmas not sustainable on Earth? Where are they possible?

They require lots of energy to maintain which we cannot constantly consistently create/input on Earth. However, plasmas are theoretically possible in a closed container.

What happens to the pressure of a gas when you heat it in a container? What else can be changed to elicit the same results?

Heating gases gives them more energy than the container can bind, so the pressure increases. Decreasing the container volume also increases the pressure.

Collisions and the Transfer of Energy Between Particles

• What happens to the energy transfer when one particle has more energy than the other?

• What is the difference between gentle and higher energy collisions?

• Can particles transfer all their potential energy?

Any time particles collide, they transfer energy. The energy amount always tries to equal out, so the particle with more energy will transfer to the particle with less energy.

Gentle collisions have less energy transfer, stronger collisions have more energy transfer. Higher energy collisions means the particles bounce farther apart.

Particles (or anything) can only transfer energy that they are currently using. For example, if it’s moving, only that specific kinetic energy is transferred. If it’s glowing, only that emitted light energy is transferred. This is visible because a faster car would cause more damage on an object than a slower car.

Heating and Cooling Curves

* What are they?

Describe phase changes relative to the energy contained in particles.

Heating Curve

• What type of process?

• Beginning of Plateaus

• End of Plateaus

• Inclines: State, Energy, IMFs

• Plateaus: Energy of Particles/System, Phase Change

__Endothermic Process__

**Beginning of Plateaus:** The first particle gains enough potential energy to change state.

**End of Plateaus:** Last particle gains enough potential energy to change state.

**First Incline:** Particles in a solid state gain potential energy, but the IMFs remain greater.

**First Plateau:** Particles are continuously absorbing and transferring energy. Potential energy of the system is maintained until all particles have enough energy to push against IMFs. *Phase change, melting point.*

**Second Incline:** Particles in liquid state gain potential energy, which is about equal to the IMFs.

**Second Plateau:** The same as the first plateau, except the particles are gaining energy to __overcome__ IMFs. Particles must also overcome atmospheric pressure to boil. Less pressure (high altitude) means a lower boiling point. *Phase change, boiling point.*

**Third Incline:** Particles in gas state gain potential energy, overcoming IMFs.

What is the difference between vaporization and boiling?

Vaporization includes every way that a liquid turns into a solid, whereas boiling specifically involves heat.

Absolute Zero

* What is it?

A theoretical temperature at 0K or -273°C where all matter has zero kinetic energy, even electrons.

Units for Potential Energy

• What are used?

• Examples

Any unit used for energy can be used for potential energy.

Examples: Joules (J) measures energy and degrees (°) measures heat energy. Joules is a more objective unit, it can be spent as movement, heat, light, etc. It is the most universally used unit and is easy to convert.

Potential Energy

* What is it?

What a particle has stored to use to do work at any given time.

The Difference Between Solubility as a Chemical and Physical Change

* Example

__Example__

Two separate substances are dissolved in two separate containers of water. This is solubility as a physical change.

Those two mixtures are then combined and a solid is formed. This is solubility as a chemical change.

Cooling Curves

• What type of process?

• Beginning of Plateaus

• End of Plateaus

• Declines: State, Energy

• Plateaus: Energy of Particles/System, Phase Change

• Open or closed container?

__Exothermic Process__

**Beginning of Plateaus:** The first particle loses enough potential energy to change state.

**End of Plateaus:** Last particle loses enough potential energy to change state.

**First Decline:** Particles are gaseous, losing potential energy.

**First Plateau:** Particles have lost enough potential energy to change state from gas to liquid. *Phase change, condensation.*

**Second Decline:** Particles are liquid, losing potential energy.

**Second Plateau:** Particles have lost enough potential energy to change state from a liquid to a solid. *Phase change, solidification/freezing point.*

**Third Decline:** Particles are solid, losing potential energy.

Note: Since a cooling curve represents a loss of energy, it must be in an open container.

Kelvin/Celsius Conversion Formula

Kelvin - 273 = Celsius

Why do some chemical reactions require energy input while others seem like they do not?

All chemical reactions require an energy input, but some reactions can move forward with what is available in the room (room temperature, lighting, etc.) which makes it seem like there was no input required.

Exothermic

• What does it the word mean?

• What is the process?

• Example

• What would you feel if you touched it?

Means “exiting heat.”

Releases more energy than it takes in (net total), pushes energy out (cooling).

Example: decomposition reaction.

These types of reactions would feel hot if you touched them.

Endothermic

• What does it the word mean?

• What is the process?

• Example

• What would you feel if you touched it?

Means “entering heat.”

Takes in more energy than it gives out (net total) (heating).

Example: synthesis reaction.

These types of reactions would feel cold if you touched them.

Cold

• What is it?

• What is it comparable to?

• How can something become cold?

The absence of heat. Something cannot become more cold, it can only become less hot. Similar to how dark is the absence of light energy.

Something can become cold from losing energy in any way (kinetic spending, radiating light or heat, multiple at the same time, etc.)

Dynamic Equilibrium

• What does the word mean?

• When, how, and where does it form?

• Example

• What happens during realistic phase changes?

• Is something held in a closed container at a non-phase change point still a dynamic equilibrium?

• Is a heating/cooling curve a dynamic equilibrium? Why or why not?

Means “moving” and “equal.”

Forms at a phase change in a closed container. Only certain chemical reactions can form this in an open container.

Occurs when there is no observable/visible net change, but there is a 1:1 ratio exchange of particles consistently occurring.

Example

A closed container holding 100mL of H2O at 100°C.

• As one particle condenses, another evaporates.

• Happens through kinetic energy transfer, so it must be initiated by the gas particle because it has more energy than the liquid particle.

• Gas particles move around and collide with the liquid particles on the surface. It transfers its energy to that liquid particle and turns into a liquid itself, while the liquid particle receives enough energy that it turns into a gas.

Note: In reality, phase changes all occur in open systems and particles can gain or lose energy in more ways than just the kinetic transfer of collisions.

Note: Holding something in a closed container at a temperature that is not a phase change point is not one of these. There is still particle exchange (less) but not the right 1:1 ratio.

Note: A heating or cooling curve is not one of these! There is no 1:1 particle exchange and there is a visible/observable change (phase change).

Elastic Collisions

• What is it?

• Why does it happen?

• Where does it exist?

When particles collide but keep their energy balance (no loss or gain) because they both had the exact same amount of energy.

This does not really exist. It only happens in a closed container.

Inelastic Collisions

• What is it?

• Why does it happen?

• Where does it exist?

When particles collide and there is an unequal energy transfer (more transfers to less).

All collisions in real life are this type.

Evaporative Cooling

• How does it happen?

• What increases the rate of evaporative cooling? What does this feel like?

• Can water evaporate in 10°C temperatures?

Sweat particles absorbs energy from multiple sources to reach 100°C worth of energy:

• Receive 37°C worth of energy from your body temperature.

• Gases from external environment collide with sweat particles and transfer kinetic energy. More wind would increase this.

• Heat and light energy from environment provides energy to sweat particles.

The sweat particles use this energy to change state to a gas. When it does this (evaporates), it takes your body heat with it, cooling you down.

Any factor that increases external energy will increase the rate of evaporation. Water can evaporate at low temperatures as long as it receives a total of 100°C worth of energy from other sources. Increasing the rate of evaporation means you lose energy quicker and feel colder.

What is the difference between boiling and evaporating?

• What particles are affected?

• What energy is required?

• Which particles must reach the phase change point?

In boiling, the phase change happens all throughout the liquid but in evaporating it happens to the surface particles only.

Both require 100°C worth of energy (for H2O), __but__ evaporating relies on multiple sources of energy (kinetic, light, heat, etc.) and boiling relies on a constant heat source/constant energy input.

In boiling, all particles must reach 100°C worth of energy to change state, but in evaporating only the surface particles must reach 100°C.

What is the difference between light and heat energy?

Both are electromagnetic energy, but visible light has higher energy than infrared (heat).

* Visible light is more to the right than infrared light on the electromagnetic spectrum so it has a shorter wavelength and a higher frequency.

Pot of Water Boiling Example

• Energy Source

• What happens to the first gas bubbles that start to rise? Why?

• What happens to the energy of the particles when the end of the plateau is reached?

• How does external energy input affect boiling (light, wind, etc.)?

• Why do different substances have different boiling points?

The heat/energy source is below the pot of water.

Gas bubbles start to rise but do not make it to the surface of the water because the heating curve is at a plateau, and the particles are colliding with each other and distributing energy.

When the end of the plateau is reached, each water molecule takes the energy it consumed and used to change state with it when in leaves the system.

External energy input (other than the stove) does not affect the boiling rate because it only affects the surface particles. This would increase the rate of evaporation instead.

Different boiling points of substances are caused by different IMFs. This is observable because every room can have solids, liquids, and gases present with the same external energy conditions.

Humidity

• What does it mean?

• What is relative humidity?

• What is the difference between hotter and colder air?

• How does humidity affect evaporation?

More humid air means there is more water in the air.

The relative humidity is the percentage of water that the air is holding out of the total 100% that is has the potential to hold.

Hotter air can hold more water, colder air holds less water.

The air becomes saturated with water molecules when it is humid, so more humid air means less evaporation occurs because there is less space for them in the air.

Determining Melting and Boiling Points of Compounds

• Physical/Chemical Property/Change determined by what?

• Different Types of Bonds (Strength, IMFs, Examples)

• Molecule Size (Example)

Melting and boiling points are physical properties that are determined by the strength of the bonds holding atoms together in compounds. The weaker the bonds, the lower the melting point.

• True covalent compounds are the weakest (ex. O2), they have the lowest IMFs.

• Polar covalent compounds are in the middle (ex. any covalent with 2+ different types of atoms like C2H4).

• Ionic compounds are the strongest (ex. NaCl), they have the highest IMFs.

Note: Within a type of compound, larger molecules have higher melting points because they have more capacity to absorb heat, which comes from having a larger number of bonds (ex. C10H22 has a higher melting point than C2H4).

Do all particles in a container have the same amount of energy?

• When is it yes? Why?

• When is it no? Why?

In a closed container during a phase change, yes (theoretically as a system, like a plateau on a heating or cooling curve).

In all other cases, no. The particles would be gaining energy from the external environment and losing energy to the sides of the container, so they would not have the same energy.

True Covalent Bonds vs. Polar Covalent Bonds

• Examples of Each

• “Pull” Strength

• How easy is it to break the bond?

Polar bonds are covalent bonds that do not share their electrons exactly equally, resulting in a partial change on each atom (like poles in a magnetic field).

True Covalent Bonds: H2, O2, F2, Br2, I2, N2, Cl2

• Both atoms are the same so they pull on the shared electron(s) with equal strength.

• Equal strength makes them the easiest bonds to break with the lowest melting point.

Polar Covalent Bonds: H2O, C2H4, C5H10, etc.

• Different types of atoms are bonded together meaning the electron(s) are slightly closer to one of the atoms.

• Each type of atom pulls with a slightly different strength, resulting in a partial charge.

• This makes them stronger than true covalent bonds but still weaker than ionic bonds.

How do we know that true covalent bonds have the lowest melting points while ionic bonds have the highest?

• Example of Covalent

• Example of Ionic

The atmosphere is made up of true covalent bonds (N2 and O2) and even when it gets really cold outside the atmosphere does not suddenly turn to liquid.

Iron railings (ionic bonds) do not start to melt in the summer heat.

Why is water (H2O) special?

• What two things combined make it special?

• How is this different than other substances on Earth?

• What is density?

When water’s polarity/bonds are combined with the organization of the atoms within the molecule, something special happens.

Most substances on Earth’s solid form is denser than its liquid form (solid gold with sink to the bottom when placed in liquid gold), but solid water (ice) floats in liquid water because the way its bonds freeze causes diamond crystalline structures that hold pockets of air and make it less dense that its liquid form.

Liquid water is the universal solvent. All ionic substances dissociate in water because of water’s polarity. Water can dissolve mostly anything. For ionic compounds, the water molecules surround the cation with the slightly negative oxygens and the anion with slightly positive hydrogens, effectively separating the cation and anion (dissociation, homogenous mixture). For covalent compounds, water separates them into their individual molecules (heterogenous mixture).

Note: Density is __not__ the same as weight. When something is less dense, that means it has more air between its particles.

How is water represented in a diagram?

• Description of Diagram

• What kind of bonds are formed?

• Why does it look that way?

• What does this structure create?

• What does an ionic compound dissolved in water look like?

The oxygen is drawn as the largest circle at the top, with two smaller hydrogen circles below it, connected by bonds that form a triangle that is not connected at the bottom.

Hydrogen bonds form between the hydrogens of one molecule and the oxygens of another, strengthening the overall structure of the solid state of water.

Both hydrogens have a partial positive charge (so they repel, which is what creates the angular bond) and the oxygen has a partial negative charge.

These water molecules line up with the oxygens lining up with two other molecule’s hydrogens when its frozen, which is what creates the diamond crystalline structures.

If NaCl is dissolved in water, both the sodium and chloride ions would be separated by whole water molecules with the oxygens close to the sodium ion and the hydrogens close to the chloride because of their charges. Remember that this would occur in three dimensions.

Why do you put ice cubes in a warm drink?

• Less Detail

• More Detail

The heat difference acts on a gradient: the heat from the drink moves toward the less heated ice cube, which cools the drink.

Liquid particles collide with solid particles and transfer energy so that the ice melts and keeps that energy, lowering the overall temperature of the drink.

The 5 Points of the Kinetic Molecular Theory

Assumes that ideal gas molecules:

1. Are constantly moving randomly.

2. Have negligible volume.

3. Have negligible IMFs meaning no attraction between the molecules and they cannot change phase.

4. Undergo perfect elastic collisions only/do not exert force on each other.

5. Have an average kinetic energy proportional to the ideal gas’s absolute temperature (measured from 0K).

Types of Phase Changes

* Names and Type (liquid to solid, etc.)

Melting: Solid to Liquid

Vaporization: Liquid to Gas

• Boiling

• Evaporating

Freezing/Solidification: Liquid to Solid

Sublimation: Solid to Gas

Deposition: Gas to Solid

Condensation: Gas to Liquid

Hydrogen Bonds

• What is it?

• What is required for it to form?

• What is another name for it?

A weak charged based attraction between a hydrogen atom and another polar molecule (not the molecule that the hydrogen is a part of).

Polar covalent compounds that contain hydrogens form these bonds between molecules.

These are also called H-bonds.

Which is stronger: the attraction that forms a hydrogen bond in water or the attraction between the ion and water molecules when an ionic compound is dissolved in water? Why?

Ionic substances dissolved in water form stronger attractions than hydrogen bonds.

This is because in a hydrogen bond, the attraction is between two partial charges, and when an ionic substance is dissolved in water the attraction is between a full change and a partial change. Since attraction strength increases with charge, the dissolved ionic compound forms a stronger attraction with the water molecules.

Pressure

• What is it?

• Physical/Chemical Property/Change?

• How and when did pressure experiments start?

The compressive force (not energy) exerted on or by a substance. This is a physical property.

People knew that gases exerted pressure but could not explain it until the 1800s when previous theories of matter (water, earth, fire, air) was challenged.

Experiments began with measuring atmospheric pressure using a barometer.

What is a barometer and how does it work?

A barometer is a shallow pan filled with liquid mercury with an upside down glass tube marked with millimeter increments on it.

A barometer works by sucking all the air out of the glass tube before flipping it so that when the atmospheric pressure compressed uniformly on shallow pan, the liquid mercury rose up in the glass tube until the the pressure in the tube was equal to the atmospheric pressure. This would give the atmospheric pressure measurement in mmHg.

What standard measurements were created from the use of a barometer? Which is more useful for labs and why? What causes the difference in measurements?

Standard Temperature and Pressure (STP) at sea level and 0°C = 760mmHg = 760 torre = 101.3kPa = 1atm.

Standard Ambient Temperature and Pressure (SATP) slightly above sea level and 25°C = 100.0kPa.

SATP is more useful for labs because rarely do labs have the condition of 0°C and sea level.

SATP is different because the temperature increase raises the pressure and the elevation lowers the pressure by different amounts.

Note: These values are never counted with respect to significant figures, just as periodic table values never count either, because they are already known, previously rounded values.

Why does water (H2O) have a much higher melting point than methane (CH4) even though they have a similar size?

Water is more polar than methane because oxygen is father from hydrogen on the periodic table than carbon.

This means that the electrons sit closer to the oxygen in water than methane’s electrons sit to the carbon, meaning that water’s atoms are slightly more changed than methane’s leading to stronger polar covalent and hydrogen bonds.

Stronger bonds require more energy to break, hence the higher melting point.

Why does hydrogen create special bonds?

Hydrogen is a nonmetal that behaves as a metal because it has one electron that it really wants to get rid of. It pushes its electron (“does the most”) farthest away which creates “strong” partial charges and forms H-bonds.

Why does water have a high boiling point compared to other substances with similar masses/bond types?

Water has more forces it must overcome to change phase (IMFs, H-bonds, atmospheric pressure) which requires more energy.

Why do ammonia (NH3) and water (H2O) have similar melting points?

They have similar polarity (right next to each other on the periodic table and both bonded to hydrogen).

Why does the temperature of the system remain constant when a liquid is boiled until all the liquid is gone (closed and open container)?

In an open container, evaporation takes heat from the top while the bottom of the system is heated, so there is no increase.

In a closed container, it is because the particles are absorbing and transferring energy between each other?

Why is it more difficult for a liquid being boiled to overcome atmospheric pressure than its IMFs?

The particles are transferring energy to each other and energy is leaving the system at the same time so they cannot gain enough energy to overcome the atmospheric pressure even though they have enough energy to overcome the IMFs from the heat source.

What are some examples in which humans take advantage of the unique IMFs of water?

Cooking (oven, pressure cooker, boiling), travel (ice roads), and recreation (snowmobiling).

What is a monometer? What value do you use for atmospheric pressure? Why are open ended monometers less efficient? Which end has a higher pressure in a monometer?

A device used to measure the pressure of a trapped gas against atmospheric pressure (open ended) or another known pressure value (closed ended).

If the atmospheric pressure is given, that value is used in the calculation. However, if no pressure value is given then assume it is STP = 760mmHg.

Open ended monometers are less efficient because you would need to find the atmospheric pressure with a barometer before you could use the monometer.

The lower end of a monometer has a higher pressure (it is pushing harder).

Note: Be able to do math associated with monometers. Always keep units and 5 significant decimals during calculations until rounding your final answer to the appropriate amount of significant figures.

Note: For barometer calculations, the ruler is always in cm and the pressure measurements are always in mmHg unless stated otherwise.

Do conversion factors come from standard or calculated values?

Standard values only, always.

Explain the math associated with a monometer.

1. Decide whether the atmospheric pressure or the gas’ unknown pressure is greater.

2. Decide the difference in cm on the ruler between the levels of mercury.

3. Convert the cm difference to mm.

4. Convert the atmospheric pressure value to mmHg if needed.

5. Add (gas pressure is greater) or subtract (atmospheric pressure is greater) the calculated mmHg pressure difference to the atmospheric pressure value (also in mmHg).

6. Convert the calculated gas pressure to different units if needed.

7. Round the calculated gas pressure to the appropriate amount of significant figures.

Explain the can demonstration for pressure.

1. The can is empty and filled with air. The atmospheric pressure inside the can is equal to the atmospheric pressure outside the can.

2. Water is added to the can and boils/evaporates. The water vapor displaces the air and pushes it outside of the can. Now the vapor pressure inside the can is equal to the atmospheric pressure outside the can.

3. The can is flipped over into cold water which condenses (gas to liquid) the water vapor and creates a vacuum. With nothing left to opposed the atmospheric pressure outside the can, the can is crushed by the atmospheric pressure. The vapor pressure inside the can is much smaller than the atmospheric pressure outside the can.

Electronegativity

• What is it and how can you tell?

• What does this mean for different types of bonds and their strength?

Elements on the right side of the periodic table really want to gain electrons (high electronegativity, more pull on electrons) and elements on the left really want to get rid of electrons (low electronegativity).

A greater difference between electronegativity values in elements in compounds makes a stronger bond. This is why ionic compounds are the strongest.

True covalent compounds have an electronegativity difference of zero between atoms (because it’s the same atom multiple times).

How do IMFs affect the boiling point of a substance?

Higher IMFs means that the boiling point will be higher because more energy is needed to overcome them.

What two forces must a liquid particle overcome to change phase through vaporization?

IMFs and atmospheric pressure.

Significant Figures

• Counting

• Final Answer (Multiplication and Division)

If there is a __decimal__ in the number, count from __left to right__ starting at the first non-zero.

If there is __no decimal__ in the number, count from __right to left__ starting at the first non-zero.

Keep the least number of significant figures in your answer of the ones you used for your calculations, excluding known values. For example: 5.0 x 2.00 = 10. because 5.0 has 2 significant figures and 2.00 has 3, so your answer (10) must have 2 because that is the least amount.

Scientific Notation

Put the decimal after the first significant digit and keep as many after as you need.

A positive exponent on the 10 means it represents a very large number, and a negative exponent means a very small number.

The exponent on the 10 represents the amount of times the decimal is moved.

Example: 3.99x10^6 = 3,990,000 with 3 significant figures.

Example: 3.9x10^-6 = 0.0000039 with 2 significant figures.

Vapour Pressure Curves

• Characteristics of Graph

• Difference Between Substances

• Atmospheric Pressure

• How can you tell which substance has the highest vapour pressure, highest IMFs, and highest evaporation rate?

The dotted horizontal line represents atmospheric pressure.

Each curve represents the vapour pressure of a substance relative to the temperature. Substances have different curves because they require different heat energy and have different IMFs.

Parts of the curve below the atmospheric pressure line is in a liquid state, and above is a gaseous state/boiling.

The part where the substance’s curve and the atmospheric pressure line intersects is the substance’s normal boiling point, therefore it changes if atmospheric pressure were to change. For example, if atmospheric pressure were to lower, the boiling point of each substance would also lower. At certain temperatures and atmospheric pressures, only some substances will boil.

The sharpest curved substance is the one with the highest vapour pressure. This can also be observed by looking at the temperature of 0°C and seeing which substance has the highest pressure value.

From this, the substance with the highest vapour pressure also has the quickest evaporation rate and the weakest IMFs.

What is another name for alcohol?

Ethanol

Vapour Pressure

• What is it?

• What creates it?

• Physical/Chemical Property/Change

• How is vapour pressure related to temperature?

Vapour pressure is the pressure exerted by a gas, it is a physical property.

It is created when liquid particles gain enough energy to turn into gas in a confined space where the particles become cramped, with too much energy (pressure) in what the space can hold.

Therefore, adding more energy to gas particles will also increase vapour pressure.

Increasing the temperature of particles increases the vapour pressure, since temperature is a measurement of energy and vapour pressure increases with energy.

What is the density of water?

1 g/mL

What is the boiling point of a substance in relation to vapour pressure and atmospheric pressure? What is one way you could decrease the boiling point of a substance?

The boiling point of a substance begins when the vapour pressure is equal to the atmospheric pressure (can observe on a vapour pressure curve), which the vapour pressure then exceeds as it changes to a gas.

You could bring a substance to a higher elevation to decrease its boiling point. This works because there is less atmospheric pressure the higher you go, so it would take less energy for the vapour pressure to exceed it.

Specific Heat Capacity

• Physical/Chemical Property/Change

• What does it measure and in what conditions?

• What is its relation to IMFs?

• What is it used for?

• What is the specific heat capacity of water?

A physical property that measures the amount of energy (in joules, J) required to raise 1 gram (g) of a substance by 1°C in a closed container.

The more energy needed (the higher the specific heat capacity) the higher the IMFs of the substance are.

Used to measure exactly how much energy input is needed for a state change or for a specific reaction to work most efficiently.

The specific heat capacity of water is 4.18 J/g°C.

Specific Heat Capacity Formula

• What is it?

• What do all the variables mean?

• Which variables can be negative?

• What are all the units?

• What is the formula for a temperature interval?

q = mcΔT

q is the amount of energy required. It is measure in joules (J) and can be positive or negative (1KJ = 1000J). A negative energy amount means that the substance is losing energy, while positive means it is gaining.

m is the mass. It is measured in grams (g) and can only be positive (1Kg = 1000g).

ΔT is the change in temperature. It is measured in degrees Celsius (°C) and can be positive or negative (0K = -273°C). A negative temperature change means the substance is cooling, while positive means it is heating up.

c is the specific heat capacity (the physical property). It is measured in joules per grams degrees Celsius (J/g°C) and can only be positive.

The formula for a temperature interval is the final temperature minus the initial temperature: ΔT = Tf - Ti.

Specific Heat Capacity Problem For One Substance

1. List out known variables and determine the missing one.

2. If needed, find the temperature interval using ΔT = Tf - Ti.

3. Make sure that all the values you have are in the correct units. If not, convert them.

4. Write out the formula.

5. Rearrange the formula so that the unknown variable is isolated.

6. Plug in all the values, remembering to use units and 5 significant decimal places where applicable.

7. Write out answer to 5 significant decimal places.

8. Rewrite final answer using significant figures.

What is a way that “five hundred grams of water” could be written?

500gH2O

Specific Heat Capacity Problem For Two Substances

1. List out known variables (using subscripts to indicate which substance they belong to) and determine the missing ones.

2. If needed, find the temperature intervals using ΔT = Tf - Ti. Usually, the initial temperature of each substance and the final temperature of the system will be given. In this case, the final system temperature is also the final temperature for each individual substance.

3. Make sure that all the values you have are in the correct units. If not, convert them.

4. Write out the formulas with the subscripts, equating the two equal variables (which is frequently q).

5. Rearrange the formula so that the unknown variable is isolated.

6. Plug in all the values, remembering to use units and 5 significant decimal places where applicable.

7. Write out answer to 5 significant decimal places.

8. If the answer is negative if finding mass or specific heat capacity, take the absolute value of the answer, since neither can be negative.

9. Rewrite final answer using significant figures.