Physical Properties of Matter

• Matter: Anything that has _____________ and ____________________.

– Mass: Amount of ______________ in an object.

Instrument:

– Volume: Amount of ______________ an object takes up.

Instrument:

Types of Matter

• Types of Matter:

1.) Pure Substances

2.) Mixtures

Link – Matching Game

Pure Substances

1.) Pure Substances – All samples have the same composition and properties.

There are 2 types of Pure Substances:

1.) Element 2.) Compound

Pure Substances

a.) Element: A substance that cannot be

chemically broken down.

Eg. Periodic Table (Li, C, Be, . . . all)

Pure Substances

b.) Compound: A substance that can be chemically broken down.

Eg.

1.) NH3 ____________ +_____________

2.) CO2 ____________ + _____________

Pure Substances - Compounds

2 types of Compounds:

– i.) Binary – A compound made from _____

different elements.

Eg. NH3, CO2, H2O

– ii.) Tertiary / Ternary – A compound made of

____ or more elements.

Eg. C6H12O6 (glucose), NaOH, KClO3

Pure Substances - Compounds

** The composition (or, % composition) of

compounds can never vary! **

(Composition is Constant).

Eg. H2O -------------- ratio: ____ mol H : ____ mol O

%H and %O must be the same in all

samples of H2O, regardless of mass!

Ex.1:2 Samples of a Compound . . .

Sample 1 Sample 2

60% X40% Y

14g X % X = ? 6g Y

Are Samples 1 and 2 the Same or Different Compound??? Explain.

Ex.2:

3 Samples of Compounds . . . Are they the SAME or DIFFERENT??? . . . Method 1: %Composition

Sample 1 Sample 2 Sample 3

X 6g 30% 1.5g

Y 12g 60% 3.0g

Z 2g 10% 0.5g

Same or Different? Explain.

Ex.2: (Use data from previous problem)

3 Samples of Compounds . . . Are they the SAME or DIFFERENT??? . . . Method 2: Ratios (X : Y : Z)

Sample 1 Sample 2 Sample 3

Original Ratio

Reduce

Same or Different? Explain.

Ex.3:

3 Samples of Compounds . . . Are they the SAME or DIFFERENT??? . . . Method 1: %Composition

Sample 1 Sample 2 Sample 3

X 12g 50% 3.5g

Y 4g 30% 2.5g

Z 2g 20% 1.0g

Same or Different? Explain.

Ex.3: (Use data from previous problem)

3 Samples of Compounds . . . Are they the SAME or DIFFERENT??? . . . Method 2: Ratios (X : Y : Z)

Sample 1 Sample 2 Sample 3

Original Ratio

Reduce

Same or Different? Explain.

Mixtures

2.) Mixtures: 2 or more substances

(“ingredients”) physically combined.

2 Types:

a.) Homogeneous Mixtures

b.) Heterogeneous Mixtures

Homogeneous Mixturesa.) Homogeneous mixtures: uniform

composition and properties throughout

the sample.

Eg. glass, metal alloys, solutions, pure airair = (21% O2, 75% N2, 4%

Other)

** Must have same (equal) amounts of substance

throughout the entire mixture. **

Homogeneous Mixtures

** Solutions are Always homogeneous mixtures!!! **

Solution: A homogenous mixture consisting of a solute dissolved in a solvent.

Eg. 1.) NaCl(aq) (aq) = ______________

2.) Other solutions?

Link-Salt dissolving Salt dissolving(2)

Heterogeneous Mixtures

Heterogeneous mixtures: No uniform composition and properties throughout the sample.

Eg. Soil, fog, cereal, salad, air (polluted).

Pure Substance or Mixture?

Item Type of Matter

Iron

Orange juice

Magnesium chlorate

Vinegar

Hydrogen

Classroom

Paint (in can)

Gatorade

Compare / Contrast

Compounds Elements Mixtures

Breaks down chemically

Pure Substance

Composition varies

Properties always the same

Homogeneous

• Task: Draw 3 boxes.

In box #1, illustrate a Mixture of Copper (Cu) and Sulfur (S). In box #2, illustrate a Compound of Cu and S.

In box #3, illustrate a different Compound of Cu and S.

KEY: Cu = shaded circle S = unshaded circle

Separating Mixtures• There are six (6) main methods used to

separate various types of mixtures:

1. Chromatography

2. Filtration

3. Evaporation

4. Magnetism

5. Separatory Funnel

6. Distillation

Separating Mixtures

• See handout on Separating Mixtures

Phases of Matter

1.) Solids 2.) Liquids 3.) Gases

Comparing Solids, Liquids, and Gases

Phase changes

Properties of Phases

Type Shape VolumeMotion

of Particles

Arrangement

of Particles

Entropy Symbol

SOLID

LIQUID

GAS

KE = Kinetic Energy (energy in motion)

PE = Potential Energy (stored energy)

Entropy: The measure of randomness of a system

Low entropy = organized

High entropy = not organized (random)

Diagram of Each Phase

Solid Liquid Gas

Pattern No Pattern No Pattern

Phase Changes:

A.) solid liquid gas: B.) gas liquid solid:

1.) Particles must speed up 1.) Particles must slow down

2.) Particles must spread apart 2.)

3.) KE increases 3.)

4.) PE increases 4.)

• Endothermic – Energy must be absorbed.

• Exothermic – Energy must be released.

Phase Changes

Type of Change Name

Solid liquid Fusion

Liquid gas

Solid gas Sublimation

Gas liquid

Liquid solid Solidification (or _________________)

Gas solid Deposition

Examples of Exothermic phase changes:

Type of Change Name

Examples of Endothermic phase changes:

Type of Change Name

Q: Which of the following phase changes is endothermic?

1. H2O(l) H2O(s)

2. H2O(g) H2O(l)

3. H2O(l) H2O(g)

4. H2O(g) H2O(s)

Heating and Cooling Curves Link: Heating Curve Tutorial

Link: Heating Curve Link

Link: Heating Curve of Water (Step-by-Step) (Link 2)

Link: Phase Changes (Notebook)

Joules

• Joules (J) – Unit for measurement of heat

1 Kilojoule (1 kJ) = ______ J

Heating Curve Graph Questions:** The graph represents a heating curve of a

substance where heat is being added at a rate of 100 Joules per minute (100 J/min.)**

1.) In the graph, where is the substance

a.) a solid onlyb.) a liquid onlyc.) a gas onlyd.) existing both as a solid and liquide.) a liquid and a gas together

2.) What is the melting point temperature of the substance?

3.) What is the boiling point temperature of the substance?

4.) How long did the fusion of the substance take?

5.) How long did the vaporization of the substance take?

6.) How long did it take to completely melt the substance?

7.) How long did it take to completely boil the substance?

8.) How much heat was needed to melt the substance?

9.) How much heat was needed to boil the substance?

1.) In what region is the substance both a liquid and a gas?

2.) How much time does it take for the solid to melt (fuse) completely? (start at 0)

3.) Name ALL regions where KE increases:

4.) What is the change in temperature from CD?

5.) What is the phase changing process at DE called? What is it called at BC?

6.) Name the region (range) where any trace of liquid is present.

Temperature

• Temperature – Measures the Average Kinetic Energy (KE = energy in ___________).

(**Does NOT measure heat!**)

Eg. (Drawing) H20(l) molecules moving at speeds of 40mph, 20mph, and 30mph.

What is the Average?

• Thermometer – Measures Temperature

**There are three (3) different units of measurement of temperature . . .

Measuring Temperature

Unit of Measurement

Water Freezing Point

Water Boiling Point

Fahrenheit (◦F)

Celsius or Centigrade (◦C)

Kelvin (K)

Link: Comparing Fahrenheit, Celsius, and Kelvin

** Converting Celsius Kelvin and Kelvin Celsius: **

◦C = K – 273 K = _______________

Conversions

Kelvin (K) Celsius (◦C)

0K

0◦C

100◦C

100K

1.) Heat ALWAYS flows from hot cold.

** We “feel” cold when our body (loses / gains) heat.

** We “feel” hot when our body (releases / absorbs) heat.

Eg. 1.) Ice pack 2.) Heat pack / heating pad

2.) Heat continues to flow from hot cold until it reaches equilibrium (equal heat on both sides), unless interrupted.

Eg. 10◦C beaker of water + 70◦C heated stone

3.) ** Mass is important!** The more mass there is, the more heat energy it can contain.

(Draw the following)

Eg. 1. Which has more energy:

2g of H20 at 100◦C or 100g of H20 at 100◦C?

Eg. 2. Which has more energy:

10ml of H20 at 50◦C or 10ml of H20 at 20◦C?

Law of Conservation of Energy

Law of Conservation of Energy: Energy cannot be created or destroyed, but it can change forms.

**The energy lost by one object (exothermic), must be gained by another (endothermic).

ENERGY LOST = ENERGY GAINED

Law of Conservation of Energy

Eg.1: Describe what happens when you put your hand in cold water.

Hand Cold Water

1.)

2.)

1.)

2.)

Law of Conservation of Energy

Eg.2: Light Bulb

Electricity filament in light bulb Light

** a lot of energy wasted as heat it goes up and out.

Measurement of Heat Energy(**refer to Table B)

q = mCΔT = Amount of heat given off or absorbed in a reaction.

q = heat (in joules (J)) J = unit of measurement for heat

m = mass (in grams (g))

C = specific heat capacity of a substance (in J/g◦C)

ΔT = change in Temperature [t2 – t1] (in ◦C)

Δ = “Change in”

*NOTE: kJ = kilojoules. 1kJ = 1000J

(** know this conversion!)

• Calorimeter – instrument used to measure the amount of heat given off in a reaction.

** NOT a thermometer (which measures the avg. KE!)

**Set up the equation q = mCΔT, using only the units for each symbol . . .

Eg.1. How many joules are absorbed when 50.0g of

H2O are heated from 30◦C to 58◦C?

Law of Conservation of Energy

**NOTE: Heat is measured in joules (J).

The value (in J) is not a (-)-number because heat is either lost (exothermic) or gained (endothermic).

For example,1.)10g of water cools from 40◦C to 25◦C. How much heat was

used (lost)?

2.)10g of water warms up from 25◦C to 40◦C. How much heat was gained?

3.) What is the difference between 1.) and 2.)?

Eg.2.(HW) p.51, #30-35.

Eg.3. When 2.00g of an unknown sample of a solid absorbed 1000. J of energy, its temperature increased from 30.0◦C to 48.0◦C. What is the specific heat capacity (C) of the solid?

Eg.4. When 5.00g of water at 25.0◦C absorbs 900. J of energy, what will its final temperature be?

*** q = mCΔT does not apply during phase changes, because there is NO change in temperature (no ΔT).***

**Refer to Heating Curve for phase change illustration.

Calculating Heat Energy During Phase Changes:** See Table B **

1.) Heat of Fusion (melting): The amount of heat it takes to melt 1g of a substance at its melting point.

Heat of fusion = J/g

q = (m)(heat of fusion) or, q = mHF

Heat of Fusion

• Eg.1. How many joules are required to melt 460g of ice at 0◦C?

Calculating Heat Energy During Phase Changes:** See Table B **

2.) Heat of Vaporization (boiling): The amount of heat it takes to boil 1g of a substance at its boiling pt.

Heat of vaporization = J/g

q = (m)(heat of vaporization) or, q = mHV

Eg.1. How many joules are required to boil 460g of water at 100◦C?

Eg.2. Link – Calculate Heat Energy Worksheet

Eg.3. p.51 # 36,37

Activity:

1.) Draw a Heating Curve.

2.) Label the phase or phase change for each region.

3.) Label each region with the appropriate formula for calculating heat (q=mCΔT; q=mHF; q=mHV)