## Matter

Substances - matter that always has the same composition and properties. Compounds and elements are examples of substances.

Elements - The basic unit of matter. Elements cannot be decomposed by any ordinary chemical means. A sample of an element is composed of atoms with the same atomic number.

Compounds - substances that are formed by two or more different elements and can be decomposed by chemical means. Their properties are usually different from the elements they are composed of.

Mixtures - A mixture of two or substances.

Homogeneous Mixtures - A mixture that is the same throughout. Also known as a solution. Salt water (NaCl(aq) is an example of a homogeneous solution).

Heterogeneous Mixtures - A miture which is different throughout. Soil and sand are examples of hetergeneous mixtures.

## Energy

Energy - the ability to do work.

Forms of Energy

Kinetic - energy of motion

Potential - energy due to position

Chemical - energy associated with a chemical change

Heat - energy associated with the temperature of a body or systems of bodies

Nuclear - energy due to the changes in mass of the nucleus

Law of Conservation of Energy - Energy can never be created or destroyed. Energy can only be changed from one form to another.

Exothermic Reactions - releases heat into its surroundings. The products have a less potential energy than the reactants. ∆H is negative.

Endothermic Reactions - absorbs heat from its surroundings. The products have a greater potential energy than the reactants. ∆H is positive.

Temperature - a measure of the average kinetic energy of a substance.

Thermometer - You need two fixed points (usually the freezing point and boiling point of water) to make a thermometer.

Celsius - Freezing point of water is 0°C. The boiling point of water is 100°C.

Kelvin - Freezing point of water is 273K. The boiling point of water is 373K.

K = °C + 273

In Chemistry we measure heat energy by using Joules.

q = mC∆T

q = Heat Energy measured in Joules (J).

m = Mass of H2O measured in grams (g).

C = Specific Heat of H2O = 4.2 J ⁄ g · °C

∆T = Change in temperature.

Sample Problem

Remember, there are 1000 Joules in 1 KiloJoule!

How many KiloJoules are required to change the temperature of 100g of water from 30°C to 87°C?

Change in temp= 87°C − 30°C = 57°C

Specific Heat of water = 4.2 J ⁄ g · °C

mass = 100 g

q = 100g × 4.2 J ⁄ g · °C × 57°C

q = 2.4 × 101 KJ

## Phases of matter

Matter exists in three phases solid, liquid and gas.

Sublimation - When a substance changes from a solid directly to a gas. Examples - CO2 and I2.

Diagram - heating curve

From A to B the substance is in the solid form and it kinetic energy increases, while its potential energy remains the same.

From B to C the substance is changing from a solid to a liquid at a constant temperature. This is the only time when solid and liquid phases coexist (solid - liquis equilibrium). The potential energy of the substance is increasing, while its kinetic energy remains the same.

From C to D the substance is in the liquid phase and its kinetic energy is increasing, while its potential energy remains the same.

From D to E the substance is changing from a liquid to a gas at a constant temperature. The potetial energy is increasing, while its kinetic energy remains the same. This is the only time when the liquid and gas phase coexist (liquid - gas equilibrium).

From E to F the substance is in the gas phase and its kinetic energy is increasing, while its potential energy remains the same.

## Gases

Gas Laws

Boyle's Law - as pressure increases volume decreases (an inverse relationship).

Diagram - Graph of Boyle's law

P1V1 = P2V2

Charles's Law- as temperature increases volume increases (a direct relationship).

Diagram - graph of charles's law

V1 ⁄ T1 = V2 ⁄ T2

Combined Gas Law

P1V1 ⁄ T1 = P2V2 ⁄ T2

Standard Temperature and Pressure (STP) - 273K and 101.3kPa or 1 atm or 760 mmHg or 760 torr.

At STP 1 mole of any gas occupies 22.4L.

Kinetic Theory of Gases

1. All gases move in a continuous random motion in a straight line path until they are deflected by some force.
2. Collisions cause a transfer of energy but the overall energy of the system remains the same.
3. The volume of gas particles is negligible compared to the volume of space in which they move.
4. No forces of attraction are considered to exist between gas particles. This is called an ideal gas. A real gas would deviate from this law. The most ideal gas is the gas with the least mass. The gases which act most like an ideal gas are Hydrogen and Helium.

## Liquids

Vapor Pressure - In a closed system the vapor exerts a pressure on the container and produces the vapor pressure above the liquid.

Boiling Point - As temperature increases so does the vapor pressure of a liquid. When the vapor pressure equals atmospheric pressure the liquid begins to boil. The normal boiling point of a liquid is the temperature where the vapor pressure of the liquid equals standard pressure.

Heat of Vaporization Hv - heat energy required to change a liquid to a gas.

q = mHv

Hv of H2O = 2259J ⁄ g

Heat of fusion Hf - heat energy required to change a solid to a liquid

q = mHf

Hf of H2O = 333J ⁄ g

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This site was last updated 08/23/02

Angelo Stanco