Chemistry - Semester 1
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Chapter Summaries: Select the chapter you wish to view by clicking below. (Some summaries taken from Chemistry The Central Science, Brown)
Section 1 - Matter and Measurement
Section 2 - Atoms, Molecules, and Ions
Section 3 - Percentage Composition and Empirical Formula
Useful Links in the Understanding of Chemistry
Section
1-
Mass and Measurement
Textbook
Chapters
Chapter 1 (all sections)
Chapter 2 (all sections)
Objectives
Express any number in
exponential
notation. Use this notation in working calculations.
Apply the rules of
significant
figures to calculations based upon experimental measurements.
Use the
conversion
factor approach to convert measured or calculated quantities from one unit to
another.
Use an algebraic equation to
solve for an unknown quantity, given or having calculated all the other
quantities in the equation.
Distinguish between
substance
and mixture,
mixture
and solution, and
element
and compound.
Describe the separation of
mixtures into pure substances and the properties of those substances
upon which the separation depends.
Use the
Periodic Table to locate and
Basic Concepts
Chemistry is one of the most basic of all the sciences. It
involves the structure and behavior of matter. Some of the topics we will
discuss are atomic structure, stoichiometry, thermochemistry, colligative
properties, chemical bonding, gas laws, acids and bases.
The quantitative description of behaviors and structure of
matter is a fundamental part of the study of chemistry. We will therefore
start our semester with a quick review of the following mathematical concepts.
Scientific notation
Significant figures
Uncertainty in measurements
Solving simultaneous equations
One of the fundamentals of measuring is that every value
must be associated with a unit. Several standardized systems of
measurement have been developed. These include the
MKS,
CGS
and British systems. There has been a consensus among scientists as
to which units should be used to report finding. These standards are
called Standard International Units. Conversion of units is a
common task when working in science. This is because many of the constant
used in our calculations have been developed using known units of measurement.
The process of changing units is called
dimensional analysis or the factor
label method (see below).
Mater can be broken into many groups according to numerous
criteria. Two of the broadest categories are
substances and
mixtures.
A substance is a type of matter that exhibits its own unique set of
chemical
and physical properties. Substances include all
elements and
compounds.
Groups or families, metals, nonmetals, and metalloids can further classify
elements. These are all easily distinguished by using the Periodic
Table.
Mixtures on the other hand are combinations of substance,
which do not have unique properties or set proportions by mass. In fact, each substance retains its
own set of properties in a mixture. Mixtures can be classified as
homogeneous
or heterogeneous.
filtration,
distillation,
chromatography,
and recrystallization are chemical techniques used to separate mixtures
into their components.
Sample Problems
Convert 12.5 km to inches.
12.5 km (1 mile / 1.609 km) (5280 ft / 1 mile) (12 in / 1 ft) = 492000 in. (3 significant figures)
Notice that each ser of parenthesis is a conversion factor where the numerator and the denominator are equal to each other and arranges such that the previous unit will cancel out
Section
2-
Atomic Structure
Textbook Chapters
Chapter 3 (all sections)
Objectives
That a chemical symbol may represent one or more atoms of an element; a molecular formula, one or more molecules of a substance.
Know the historical
development of the structure of the atom including the findings of
Democritus,
J. J. Thomson,
Robert Millikan,
Ernst Rutherford,
Neils Bohr, and
John Dalton.
Know the origins of the
Law
of Conservation of Mass, the
Law of Definite or Constant Composition,
and the
Law of Multiple Proportions.
Relate the numbers of
protons
and neutrons in an atom to its
nuclear symbol.
Relate the numbers of
electrons
and protons
to the charge on a
monatomic
ion.
Relate the charges on
anion
and cation to the formula for the ionic compound.
Relate the
atomic mass
of an element to the abundances and masses of it
isotopes.
Write the
molecular
formulas for
ionic
compounds.
Given the formula of a
compound,
determine its
molecular mass.
Relate the numbers of atoms
(or molecules) and the mass in grams of a sample of matter.
Basic Concepts
Most of you will enter general chemistry knowing that an atom is basic unit of most elements and that a molecule is the basic unit of many compounds. The historical account of how we have come to "know" these simple facts is quite interesting and exciting. This chapter will highlight some of the more significant events in the development of early atomic theory. The modern view of the atom will be explained later in the course when we deal with electron structure.
Although much
knowledge of chemistry had been found by alchemist during their experimentation,
a fundamental understanding of why or how reactions occurred was not
revealed. In order to answer these questions science must have a model of
matter to work with. One of the first models of matter developed was
that of
Democritus in the year 500 BC. He was a Greek philosopher
who is credited with stating the first atomic theory. He stated that
matter was composed of ultimate indivisible particles called an atoms.
This atomic theory fell from favor due to the teaching of another influential
person of the times, Aristotle. It took nearly 2 centuries for
atoms to be recognized again.
The
indivisibility of the atom was challenge by J. J. Thomson while
conducting
experiments with
cathode ray
tube. He is credited with
discovering the subatomic called an
electron. Thomson was
able to calculate the ratio of mass to charge of the electron by using
magnetic fields. A new atomic model of the atom was developed by Thomson
which is often referred to as the
"plum pudding" model.
The continued experimentation with cathode rays led to the discovery of canal
rays which were later identified as
protons.
The actual charge of the electron was found by Robert Millikan using his famous oil drop experiment. The value obtained (1.6 x 10-19 C) combined with Thomson's ratio allows for the calculation of the mass of the electron (9.1 x 10-31 kg).
The next
major development of atomic theory was contributed by Ernst
Rutherford.
The
Rutherford experiment showed that the majority of the volume of the atom is "empty" by
aiming alpha particles at a thin gold foil and noting the patterns of the
particles as they passed through. This gave rise to what is called the orbital
atom model where the majority of the mass of the atom is found in the nucleus
of the atom and the electrons are found a relatively large distance away.
A refinement of the orbit atom model was submitted by Neils Bohr a few years later. In trying to explain the spectral lines of the elements, Bohr proposed not only that the electrons were separated from the nucleus but that these electrons could only orbit the nucleus at specific distances containing specific energies.
Isotopes are atoms of the same element which have different masses. This is possible because of variances in the number of the third subatomic particle, the neutron. The neutron was discovered by James Chadwick in 1932. Since not all atoms of an element have the same mass, we needed a term that represents the average mass of large number of atoms. This is called the atomic mass of an element.
The counting unit of chemistry is the chemical mole. A mole is a specific number of particles be they ions, atoms, or molecules. That number is 6.022 x 1023, Avogadro's number. It is also important to realize that 1 mole of any element has a mass equivalent to that element's atomic mass expressed in grams. Likewise, 1 mole of a compound has a mass equal to its molecular mass in grams.
Sample Problems
Section
3-
Percentage Composition and
Empirical Formula
Textbook
Chapters
Chapter 7 (all sections)
Objectives
How to interpret
formulas - in particular, the meaning and use of subscripts.
Determine the
empirical formula of a compound, given the
mass percentages of the elements
or analytical data from which these can be calculated.
Determine the
molecular
formula of a compound, given the
empirical
formula and at least an
approximate
molecular mass.
Use the Periodic
Table to obtain the charges of ions formed by the
main-group
elements.
Write
the formula for an ionic compound given either the formulas of the
ions of the name of the compound.
Given the formula for a
compound, give its name.
Describe some experimental
methods of determining percent composition.
Basic Concepts
As we have seen compounds are composed of different combinations of elements. Each compound is unique in its composition and therefore the percentage by mass of each element in these compounds can easily be calculated. The percentage composition can be found both experimentally or theoretically. An experimental calculation is one that involves the use of measurements taken during a laboratory exercised. Theoretical calculations are performed form the known formula of a compound using the atomic masses of its elements. In either case, the percentage of an element is found by dividing the mass of the element by the mass of the total compound and of course multiplying by 100.
The empirical formula of a compound is the simplest ratio between the elements in that compound. This formula can be found by using experimental data, which would allow for the calculation of percentage compositions of the elements. The process involves the conversion of masses or percentage compositions into a ration of moles represented as small whole numbers. These numbers will become the subscripts of the empirical formula.
Molecular formulas are representation of the exact number of each element in a molecular unit of a compound. Finding the ratio of the molecular mass of a compound and comparing this value to the mass of the empirical formula can find the molecular formula.Sample Problems
Section 4-Stoichiometry
Textbook Chapters
Chapter 8 (all sections)
Chapter 9 (all sections)
Objectives
Write and balance simple
chemical
equations.
Given a
balanced equation,
relate the numbers of
moles of any two substances taking part in the
reaction.
Given the
balanced equation, relate the masses of any two substances taking
part in a reaction.
Given or having calculated
two or the three quantities -
concentration (M), number of moles
of solute, volume of solution - calculate the other quantity.
Given the
balanced equation
for a reaction involving species in solution, relate the
volumes or concentrations
of two reactant species.
Describe reactions in water
solution as involving
ionic,
gas
forming,
precipitation,
acid-base,
and
oxidation-reduction.
Basic Concepts
A chemical equation can be represent as a balanced equation which includes the formulas of chemicals, their states of matter, and how much of each will be reacted or produced. Equations are written so that reactants are on the left side of the yield sign and the products are on the right. States of matter are indicated at the end each formula using standard abbreviations. The law of Conservation of Mass controls the balancing of the equation. It indicates that the total number of atoms of each element on the reactant side of the equation must be equal to the total number each element found on the product side.
Sample Problems
Section
5-
Thermochemistry
Textbook
Chapters
Objectives
Use a thermochemical
equation to relate heat in a reaction to amounts (moles, grams) of products
or reactants.
Relate the enthalpy
changes (DH) for two reactions whose
equations differ only in the direction of reaction and/or in the values of
the coefficients used.
Calculate
DH for a reaction from enthalpies.
Use calorimetric data
to determine the heat flow, Q, for a reaction.
Calculate
DH for a reaction from heats of formation.
Basic Concepts
This chapter is been about energy and the first law of thermodynamics. Energy can be measured in terms of the ability to do work or transfer heat. An object may possess potential energy because of its position or because of its composition. Chemical energy is potential energy, which may be released when the object undergoes a chemical change. An object may possess kinetic energy because of its motion relative to other objects. The first law of thermodynamics, also referred to as the law of conservation of energy, states that in any change that occurs in nature, the total energy of the universe remains constant. Any process in which heat energy is lost to the surroundings is termed exothermic and any change which absorbs heat from the surroundings is termed endothermic.
Heat changes occurring at constant pressure are of special interest. The heat gained or lost by the system in a process occurring at constant pressure is termed the enthalpy change (DH). This quantity is negative for an exothermic reaction and positive for an endothermic process. The application of Hess's law shows that the overall enthalpy change in any process can be represented by the sum of many small individual steps. The use of standard heats of formation (DHfo) is extremely helpful in determining changes of enthalpy using Hess's law. In solving problems dealing with enthalpy changes it is important to remember the following points:
The enthalpy change in a reaction, DH, is directly proportional to the amount of substance that reacts or is produced.
DH for any reaction is equal in magnitude but opposite in sign to the value of DH for the reverse reaction.
The heat of formation for any element in its standard state is zero.
Heat changes can be measured in a calorimeter in which the heat evolved or absorbed in a reaction is measured by observing the change in temperature of calorimeter. The specific heat capacity of an object is the quantity of heat required to produce a unit change in temperature.
Sample Problems