Chapter 1: Introduction to Chemistry
Chemistry is the study of the composition of matter and the changes that matter undergoes. Chemistry can be subdivided as follows:
- Organic chemistry is defined as the study of all chemicals containing carbon.
- The study of chemicals that in general do not contain carbon is called inorganic chemistry.
- The study of processes that take place in living organisms is biochemistry.
- The area of study that focuses on the composition of matter is analytical chemistry.
- Physical chemistry is the area that deals with the mechanism, rate and energy transfer that occurs when matter undergoes a change.
A model is a representation of an object or event.
A hypothesis is a proposed explanation for an observation.
An experiment is a procedure that is used to test a hypothesis. When you design experiments, you deal with variables, or factors that can change. The variable that you change during an experiment is the independent variable, also called the manipulated variable. The variable that is observed during the experiment is the dependent variable, also called the responding variable.
A theory is a well-tested explanation for a broad set of observations.
A scientific law is a concise statement that summarizes the results of many observations and experiments.
Chapter 2: Matter and change
Matter is anything that has mass and occupies space.
An extensive property is a property that depends on the amount of matter in a sample. Examples of extensive properties include mass (a measure of the amount of matter the object contains) and volume (a measure of the space occupied by an object).
An intensive property is a property that depends on the type of matter in a sample, not the amount of matter. Examples include density, color and hardness.
A physical property is a property of a substance that can be observed or measured without changing the substance’s composition.
Matter that has a uniform and definite composition is called a substance. A substance can be either an element or a compound.
Three states of matter are solid, liquid and gas. A solid has a definite shape and volume. A liquid has a definite volume but it takes the shape of the container in which it is placed. A gas takes both the shape and volume of its container. A vapor is the gaseous form of a substance that is usually a liquid or a solid at room temperature.
A mixture is a physical blend of two or more components mixtures can be either heterogeneous or homogeneous. A homogeneous mixture is perfectly mixed even at a molecular level and is also called a solution. Examples of solutions include air, metal alloys such as bronze, and sugar dissolved in water. Examples of heterogeneous mixtures include blood, milk (a suspension of tiny drops of fat in water) and butter (a suspension of tiny drops of water in fat).
A phase is any part of a sample with uniform composition and properties. When oil and vinegar are mixed they form a heterogeneous mixture with two layers, or phases, each of which is a liquid.
Differences in physical properties can be used to separate mixtures. Example techniques include filtration, distillation, chromatography and many others.
Substances can be classified as elements or compounds. An element is the simplest form of matter that has a unique set of properties. A compound is a substance that contains two or more elements chemically combined in a fixed proportion. Compounds can be broken down into simpler substances by chemical means, but elements cannot.
Chemists use chemical symbols to represent elements and chemical formulas to represent compounds.
The ability of a substance to undergo a specific chemical change is called a chemical property. For example the ability to rust is a chemical property of iron.
A chemical change is a change that produces matter with a different composition than the original matter. A chemical change is also called a chemical reaction. During a chemical change the composition of matter always changes. One or more substances change into one or more new substances during a chemical reaction. A substance present at the start of the reaction is a reactant. A substance produced in the reaction is a product. In the burning of charcoal, for example, carbon and oxygen are the main reactants and carbon dioxide is the main product.
Possible clues that a chemical change has taken place include a transfer of energy, a change in color, the production of a gas, or the formation of a precipitate. A precipitate is a solid that forms and settles out of a liquid mixture.
During any chemical reaction the mass of the products is always equal to the mass of the reactants. The law of conservation of mass states that in any physical change or chemical reaction, mass is conserved. Mass is neither created nor destroyed.
Chapter 3: Scientific Measurement
In scientific notation, a given number is written as the product of two numbers: a coefficient and 10 raised to a power. For example, 3040.0 becomes 3.0400 x 103.
Accuracy is a measure of how close a measurement comes to the actual or true value of whatever is being measured. Precision is a measure of how close a series of measurements are to one another, irrespective of the actual value.
There is a difference between the accepted value, which is the correct value for the measurement based on reliable references, and the experimental value, the value measured in the lab. The difference between the experimental value and the accepted value is called the error.
The percent error of a measurement is the absolute value of the error divided by the accepted value, multiplied by 100%.
The significant figures in a measurement include all of the digits that are known, plus a last digit that is estimated. For the rules for counting significant figures, see page 67.
In general, a calculated answer cannot be more precise than the least precise measurement from which it was calculated. The calculated value must be rounded to make it consistent with the measurements from which it was calculated.
For addition and subtraction, the answer calculated should be rounded to the same number of decimal places (not digits) as the measurement with the least number of decimal places. For multiplication and division you will need to round the answer to the same number of significant figures as the measurement with the least number of significant figures.
In science we use the International System of Units (abbreviated SI for ‘Système International’ for it was invented in France). SI units include the meter (m), the kilogram (kg) and the second (s). For a table of metric prefixes, see table 3.2 on page 75.
Forces are measured in newtons (N). Weight is the force of gravity acting on an object and is therefore measured in newtons, not kilograms.
The SI unit for energy is the joule (J) but energy can also be measured in calories (C).
Scientists measure temperature either in degrees Celsius or in kelvin. To convert a temperature in degrees Celsius into kelvin, simply add 273. Zero kelvin is absolute zero, the coldest possible temperature.
Density, an extensive property, is equal to mass divided by volume, so in the SI system it has units kg/m3.
To convert measurements into different units, use a conversion factor. For example, to convert 750 dg into grams (see page 88): 750 dg x (1 g / 10 dg) = 75 g (1 dg = 1 decigram = 0.1 g)
Chapter 4: Atomic Structure
An atom is the smallest particle of an element that retains its identity in a chemical reaction. For the history of the discovery of the structure of the atom see pages 102-109 of your textbook and the corresponding photocopied notes.
The atomic number (proton number) of an element is the number of protons in the nucleus of an atom of that element. This determines what element the atom is. For example, any atom containing 6 protons MUST be carbon (see the periodic table to obtain the atomic number of each element). In a neutral atom the number of electrons must equal the number of protons so that the negative charge of the electrons exactly cancels the positive charge of the protons.
A nucleon is a nuclear particle i.e. a proton or a neutron.
Atomic mass number (nucleon number) is the number of nucleons i.e. the total number of protons and neutrons. That means that to find the number of neutrons you must subtract the atomic number from the mass number.
The modern periodic table of the elements lists the elements in order of their atomic number (proton number) and groups into columns elements with similar properties. The columns of the periodic table are called groups and the rows are called periods. For more information on the periodic table, see chapter 6 of your textbook or the photocopied sheets.
A nuclide is a distinct kind of atom or nucleus characterized by a specific number of protons and neutrons. Example of a nuclide: carbon-14 (C-14 or 614C) contains 6 protons and 8 neutrons. People often wrongly use the word 'isotope' when they should use the word 'nuclide'. For example, they might say "Carbon-14 is an isotope" but isotope means 'the same number of protons' and it is meaningless to say "Carbon-14 has the same number of protons". The same number of protons as what? It is better to say "Carbon-14 is a nuclide".
Isotopes are variations of the same element that have the same number of protons but different numbers of neutrons. For example carbon-12 and carbon-14 are isotopes because they have the same number of protons (6) but different numbers of neutrons (6 and 8 respectively). Isotopes have the same chemical properties since they are the same element.
The atomic mass unit (amu) is one twelfth the mass of a carbon-12 (C-12) atom. 1 amu is almost exactly equal to the mass of a hydrogen-1 atom, the common form of hydrogen and the lightest atom of all, so if we say that the mass of a nitrogen-14 atom is 14.003 amu then we mean it is 14.003 times heavier than the lightest atom, a hydrogen atom. The mass of a proton and the mass of a neutron are both almost exactly equal to 1 amu, while the mass of an electron is equal to 1/1840 amu. One amu is equal to 1.67 x 10-24 g.
The (relative) atomic mass of an element is the weighted average of the masses of the isotopes of that element. For example, chlorine exists as two isotopes, Cl-35 and Cl-37. To get the weighted average we have to take into account the relative abundance of the different isotopes in nature. For example, 75% of chlorine atoms in nature are Cl-35 and 25% are Cl-37. The mass of an atom of Cl-35 is almost exactly 35 amu and the mass of a Cl-37 atom is almost exactly 37. Therefore the atomic mass (the weighted average) can be found with (0.75 x 35) + (0.25 x 37) = 35.5 amu.
Atomic structure: At high school level we usually use a model of the atom in which very light, negatively charged electrons orbit a relatively heavy nucleus consisting of positively charged protons and electrically neutral neutrons. A good analogy is the solar system, where relatively light planets orbit around the heavy sun. In chemical changes, the nucleus of the atom never changes, but electrons, being on the outside, can be gained, shared or lost by the atom. If a neutral atom gains or loses one or more electrons then it becomes a charged atom, called an ion.