|Intro||Ch. 5: Chemical Formulas|
|Ch. 1: Matter||Ch. 6: Chemical Quantities & the Mole|
|Ch. 2: Measurements & Significant Figures||Ch. 7: Balancing Equations & Reaction Types|
|Ch. 3: Factor Label Method||Ch. 8: Stoichiometry|
|Ch. 4: Introduction: Atomic Stucture|
Matter is defined as anything that has mass and volume. Anything that has mass also has inertia, that is it takes force to put matter into motion and force must be used to stop its motion.
Matter is classified in a number of ways based on its chemical and physical properties. For example the state; (solid, liquid or gas) matter exists at for particular pressure and temperature are physical properties.
Matter may also be classified based on composition as homogeneous or heterogeneous mixtures or as pure substances; elements or compounds.
Mixtures contain more than one pure substance; any combination of different elements and or compounds.
Elements are pure substances that are all composed of one "kind" of atom i.e. they must be found on the periodic table of elements. Examples: Iron (Fe), Hydrogen (H), and Oxygen (O).
Compounds are pure substances that are composed of 2 or more different elements that are held together by chemical bonds i.e. they must be sharing electrons or are attracted to each other do to charges from the loss and gain of electrons. Examples: water (H2O), table salt (NaCl), sugar (C6H12O6).
There are two major kinds of mixtures heterogeneous mixtures and homogeneous mixtures.
Heterogeneous mixtures contain more than one kind of pure substance which is not uniform on a molecular or atomic level. Usually we can see the mixture is not uniform because the "clumps" of matter are large. Examples are concrete, chocolate chip cookies, soil, and salsa. Sometimes the "chunks" are so small it may appear the mixture is uniform but is not.
Example: Milk is really a Heterogeneous mixture of water, fat droplets and dissolved substances. Liquid Heterogeneous mixtures can be detected using the Tyndall effect.
Physical properties are the properties of a pure substance which can be determined or measured without changing the chemical properties of the substance, that is without changing it from one substance to another.
Examples of Physical Properties are:
Density, solubility, boiling point, condutivity, specific heat, hardness, ductility, mallability, magnetism, color, freezing point.
True mixtures can be separated into their pure substance constituents using physical properties.
Example: a mixture of iron filing and sulfur can be separated using a magnet because iron is magnetic and sulfur is not. Note this separation does not change the chemical nature of the substances.
A mixture of water, alcohol and sugar can be separated using the substances boiling points and differences in solubility. How would you seperate this mixture.
Physical properties are often very useful for the identification of a substance.
Several physical properties often create a unique profile for an individual substance.
Chemical Properties are the properties of a substance that relate to its reactivity with other substances, that is how readily it combines with other substances to make compounds or how readily a compound breaks apart to form different compounds or elements.
Examples of chemical properties are how fammable, corrosive, explosive or reactive a substance is with another chemical.
Chemical reactions often release or require heat, give of light, involve color changes, or involve the release of a gas.
Some important derived physical properties:
Derived physical properties are properties relating two or more extensive properties to create an intensive property.
Extensive properties depend on one of the following: mass, size, temperature, length or volume of a substance.
A derived physical property compares two or more of these to produce a new "constant" extensive property.
Example: Density is the ratio of mass to volume, i.e. mass/volume.
This ratio is now is contant for that substance at a constant temperature and pressure.
Example: The density of water is 1.0grams/mL at 4ºC and 1.0 atmospheres.
What this means is at 4º C and 1.0 atm. of pressure it does not matter whether you have 1.0 qts of water or 1000.0 gal. the density will be 1.0 gram/mL.
Problems with density: The most standard units of density are g/mL which is the same as g/cm3. Therefore when solving density problems our answer should have units of g/mL.
Examples: What is the density of a piece of metal with a mass of 50 grams and has a volume of 12 mL?
Since grams should be our unit on the top and mL the unit on the bottom. 50 will be in the numerator and 12 will be the devisor in the denominator. The answer is 50g/12mL= 4.12g/mL.
Some times we have the density of a substance and want to find its mass or volume. The following problems are examples of those kinds of problems: Try to find the correct answer.
What is the volume of 80.0 g of ether if the density of ether is 0.70 g/mL?
a. 56 mL c. 8.9 x 10-3mL
b. 1.1 x 102mL d. 6.0 x 102mL
What is the mass in grams of a cubic meter of balsa wood if the density of balsa wood is 0.20 g/mL?
a. 2.0 x 10-2g c. 2.0 x 103g
b. 2.0 x 10-1g d. 2.0 x 105g
Concentrated hydrochloric acid has a density of 1.19 g/mL. What is the mass, in grams, of 1.00 liter of this acid?
a. 1.19 x 103g c. 0.840 x 102g
b. 1.19 x 10-3g d. 0.040 x 10-2g
Another important derived physical property is specific heat. Specific heat tells us how much energy is required to raise the temperature of one gram of a substance one degree celsius. It is a ratio of energy to mas and change in temperature.
specific heat = energy/mass*temperature.
Again water was used as the standard for specific heat. The amount of energy that is required to raise one gram of water one degree celsius is called the calorie. Most substances have specific heats less than water, therefore less than 1cal/g*ºC.
In fact most metals have specific heats from 0.09-0.3 cal/g*ºC. The standard SI units for specific heat are joules/g*ºC however we will primarily be using cal/g*ºC.
Example problem for specific heat: A piece of gold weighing 35.0 g absorbs 185 J of heat energy when its temperature increases by 410C. What is the specific heat of gold in cal/g*ºC?
Again we think of the units we want in the answer which is cal/g*ºC and we use the information in the problem and arrange exactly the same in our calculation. In this problem we must convert joules to calories.
During chemical reactions both mass conserved. That is to say that you always end up with the same amount of matter in the products as you started with. For example if you start with 4.0 grams of hydrogen and 32.0 grams of oxygen and react them together in a chemical reaction you will end up with 36.0 grams of water.
2H2O + O2 --------------> 2H2O 4.0 +32.0 = 36.0 grams of reactants =36.0 grams of product.
The same is true for all non-nuclear reactions. If 224.0 grams of Iron react with 96.0 grams of oxygen 320 grams of iron III oxide will be formed:
4Fe + 3O2 -------------------->2Fe2O3 224.0g + 96.0 g =320 grams of reactants =320 grams of product.
The conservation of energy is not so easy to see because it is being converted from one form of energy to another. In these example cases some of the "chemical energy" , that is the potential energy stored in the bonds of the chemical compounds is being released i.e. converted into heat energy in the same way that when wood burns it reacts with oxygen in the air and chemical energy stored the wood is released as heat and light energy. The first law of Thermodynamics states that energy cannot be created nor destroyed it can only be converted from one form to another.