The Properties of Matter

Unit 2

The Properties of Matter:

Chemistry: It is the study of Matter and its Composition, Properties, and the Changes that it undergoes.

1. Matter: it is the stuff that occupies space, and has mass, and which can be changed by energy.

A. Mass: the amount of matter that the object contains. For example, a golf ball would have more mass than a tennis ball … why? UNITS of Mass: grams, kilograms, pounds

B. Volume: the amount of space that the object occupies. UNITS of Volume: liters, milliliters, gallons For example, the tennis ball has greater volume than the tennis ball … what can we say about volume and mass? Do they go hand in hand?

2. Composition: What is it matter made of? A. PURE SUBSTANCE (or Substance): contains only one kind of matter.  Element: the fundamental substances of chemistry. There are over 100 different elements listed in the Periodic Table according to their Atomic Number.

i.e. Na -> Sodium K -> Potassium Fe -> Iron

A sample of Copper consists of MANY “atoms” of Copper (Cu).

 The ATOM: the smallest particle of an element that we can identify as that element.

The Diameter of a Copper Atom, for example, is 2.70 x 10-10 m, or 27 nanometers (1 nanometer = 10-9 meters)

When we speak of atoms in a measurable sample of an element, we often use the term “mole” to describe the amount of these tiny particles:

1 mole = 6.02 x 1023 So: 1 mole Copper atoms = 6.02 x 1023 Copper atoms

If we combine Elements, we would form COMPOUNDS, which in the same way elements are made up of discreet atoms, compounds are made up of discreet MOLECULES.

 COMPOUND: made up from the combination of 2 or more elements in a precise, well-defined ratio.

Examples: Unit 2

H2O (water) -> a molecule of water is made up of two hydrogen atoms and one oxygen atom.

NaCl (common table salt) -> made up of a ratio of one sodium atom and one chlorine atom.

C6H12O6 -> glucose; the notation represents 6 carbon, 12 hydrogen, and 6 oxygen atoms

In Summary:

Large Size < ------> Small Size

Compounds (molecules) (made up of) -> (a ratio of) Elements (made up of many) -> Atoms

B. MIXTURES: is a physical blend of two or more substances. The composition of a mixture may vary.

For example, a cup of coffee can have varying ratios of coffee, to milk, to sugar, to sweetener. Other examples of mixtures include blood (varying amounts of oxygen, water, cells, etc), air (differences in O2, N2, pollutants, etc.)

Mixtures can be divided into Two Types: (1) A Heterogeneous Mixture: is one that is not uniform in composition. For example, a dinner salad consists of several components that are not evenly distributed. CAN BE PHYSICALLY SEPARATED into substances.

What are examples of other heterogeneous mixtures?

(2) A Homogeneous Mixture: is one that has completely uniform composition; its components are evenly distributed throughout the sample.

Examples: Salt Water, a silver ring. Would a dime be an example of a homogeneous mixture?

Homogeneous mixtures are VERY important in chemistry. In the liquid phase, these homogeneous mixtures are called SOLUTIONS.

A PHASE is any part of the system with uniform composition and properties.

3. Suspension: blood. Milk, orange juice. LOOKS CLOUDY, uniform. Particles will separate out (precipitate out) of mixture is left to settle. CAN BE PHYSICALLY SEPARATED into substances. Unit 2

4. Colloid: gelatin, interior of a living cell. Particles remain in solution indefinitely. Looks like a GEL. CAN BE PHYSICALLY SEPARATED into substances. Unit 2

What are the States of Matter? (use cup of water)

To take a detailed look at the states of matter, we must first look at Potential vs. Kinetic Energy:

Potential (Stored) Energy vs. Kinetic Energy (energy of motion):

How can Energy play a role in our description of States of Matter?

Let’s check out the PowerPoint entitled “Kinetic Molecular Theory” to find out. Unit 2

5. Properties and Changes that Matter undergoes: A. Physical Change: change in the physical appearance without change in the chemical composition. For instance, change of state of matter, from solid, liquid, to gas, etc. B. Physical Properties: density, Boiling Point, Melting Point, Freezing Point, Conductivity. Physical properties are characteristics that can be observed without a chemical reaction

These physical properties are INTENSIVE, which means that the property does not depend on the amount of material present.

For example, 1 liter of water or 1 gallon of water would boil at 100 degrees Farenheit.

Some Physical/Intensive Properties: Color, Odor, State, malleability, Ductility (metal can be drawn into a wire), Luster, Solubility, Viscosity, Luminosity, Hardness, Density, & Conductivity

Boiling Point: Temperature at which a liquid changes to the gaseous phase. For a PURE substance, the temperature remains constant as the liquid boils off.

Melting Point: Temperature at which a solid changes to the liquid phase. For a PURE substance, the temperature remains constant as the solid melts.

Freezing Point: Temperature at which a liquid changes to the solid substance.

(By contrast, EXTENSIVE Properties depend on the amount of material; i.e. mass, volume, energy.) Unit 2

Let’s check out this web-based interactive: http://www.chm.davidson.edu/ChemistryApplets/PhaseChanges/HeatingCurve.ht ml

All of the Physical Changes can be summarized:

C. Chemical Properties: describes how one substance reacts with another. D. Chemical Change: the substance undergoes a change in chemical composition, in a Chemical Reaction.

A chemical reaction is given by a Chemical Equation:

A + B  C + D Unit 2

Reactants “Yields” Products

Some example chemical reactions:

a. Sugar (compound of C12H22O11) -> Carbon (element) + Water (compound of H2O)

b. H2O + SO3 -> H2SO4 (sulfuric acid)

c. 2 H2O  2 H2 + O2

d. Combustion of a hydrocarbon: rapid combination of oxygen with other materials

i.e. 2 C8H18 + 25 O2  16 CO2 + 18 H2O

DO THE REACTANTS EQUAL THE PRODUCTS???

The Law of Conservation of Mass:

States that in any physical change or chemical reaction, mass is neither created nor destroyed. It is conserved … or simply stated, the mass of the reactants is equal to the mass of the products.

Let’s go back to our examples of chemical reactions to verify that mass is conserved:

a. H2O + SO3 -> H2SO4 (sulfuric acid)

Reactants: 2 H, 1 S, & 4 O Products: 2 H, 1 S, & 4 O

b. 2 C8H18 + 25 O2  16 CO2 + 18 H2O

Reactants: 16 C, 36 H, 50 O Products: 16 C, 36 H, 50 O

How do We Know that a Chemical Reaction Has Occurred?

What Are the Signs?

1. Bubbles Appear as gas is gas is produced from the reaction. 2. A Precipitate Forms as a new solid product “rains down” or forms at the bottom of the container. This precipitate (ppt.) is “insoluble” in the reaction mixture. 3. A color change occurs 4. The temperature changes. a. It takes energy to break chemical bonds b. When bonds form, energy is released, or given off. Depending on the balance of the breaking/forming of bonds, energy is given off or taken in. 5. Light is emitted. Energy that is given off in a reaction sometimes takes the form of light. (i.e. Bioluminescence of the firefly). 6. A Change in Volume occurs. Changes in density from reactants to products will cause a change in volume. Unit 2

7. A change in electrical conductivity occurs. Some reactions produce or consume ions (charged particles) in a solution. Changes in the character and concentrations of the ions will cause the reacting mixture's ability to conduct electricity to change. 8. A change in melting point or boiling point occurs. The melting or boiling point is characteristic of a compound; when the composition of a mixture changes, the melting point and boiling point also change. 9. A change in smell or taste occurs. Since many chemical reactions have poisonous reactants or products, this method of detecting chemical change isn't recommended! 10. A change in any distinctive chemical or physical property occurs. Pick a property that uniquely characterizes one of the compounds involved in the suspected reaction, and monitor it. If the property really distinguishes that compound from all the others, you'll see it change when a reaction occurs

~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~ E. Thermal Change: the change that occurs when matter absorbs or releases heat.

Remember: Temperature is a measure of the average kinetic energy of matter. Heat is the total energy of the sample; it is the form of energy that flows between two bodies when the bodies are at different temperatures; heat will flow from the hotter body to the cooler body.

Endothermic Change: occurs when matter ABSORBS energy (in=endo) Exothermic Change: occurs when matter RELEASES energy (exit=exo)

Examples: 1. Boiling of Water: heat is added to the water to increase it’s temperature until it ultimately boils … Endothermic change 2. Melting of Ice: heat is absorbed by the ice to become a liquid … Endothermic change 3. Condensation of water (gas to liquid), or freezing (liquid to solid): Water releases energy; these are Exothermic changes ______

Separation of Mixtures: how can we separate mixtures into their components?

a. Physical methods: use of a magnet to separate a mixture of sulfur and iron? Why can these be separated by a magnet? i. Separation based on physical properties; i.e. magnetic ii. Physical separation with tweezers iii. Filtration of a solid that does not dissolve in a liquid, but is suspended, or precipitated at the bottom of the flask. Unit 2

iv. Solubilities: we can separate mixtures based on solubilities. b. Distillation: separation of a liquid mixture based on the boiling points (physical, intrinsic property) of each component.

Examples: distillation of salt water to obtain purified water, or the distillation of wine to obtain pure alcohol.

Chromatography: separation based on solubility. Demo: separation of Food Coloring (M&M’s) in vinegar Materials: 1. 1-600 ml beaker 2. prepared sample of candy coating dye 3. vinegar 4. ruler 5. capillary tubes 6. chromatography paper (4 x 20 cm) 7. pencil & plastic wrap

Rf Value = distance traveled by the component / distance traveled by the solvent