stFor 1-year undergraduate students of CCCE
Xin Lu (吕鑫)
Tel: 2181600 (Office)
C (Fullerene)Carbon Nanotube60
Two hottest allotropic forms of carbon
• Some Aspects of General Chemistry ( ~ 7 times) • Some Aspects of Organic Chemistry ( ~ 4 times) • Some Aspects of Biochemistry ( ~2 times)
URL of this course: http://pcss1.xmu.edu.cn/~xinlu/courses/ce/
• **Introduction to General, Organic and Biological Chemistry by Sally
Solomon, McGraw-Hill Book Company.
• Chemical & Engineering News (Weekly), American Chemical Society.
Chapter 1 Measurements ??????????????? 3 Chapter 2 Matter and Energy ?????????????? 5 Chapter 3 Atoms ???????????????????????? 7 Chapter 4 Chemical Bonding ???????????????? 12 Chapter 5 Gases and Atmosphere ????????????? 18 Chapter 6 Liquids and Solids ???????????????? 20 Chapter 7 Solutions ?????????????????????? 23 Chapter 8 Chemical Reactions ???????????????? 28 Chapter 9 Acids and Bases ?????????????????? 32 Chapter 10 Alkanes ??????????????????????? 35 Chapter 11 Alkenes and Alkynes ??????????????? 38 Chapter 12 Benzenes and The Aromatic Hydrocarbons ??? 40 Chapter 13 Alcohols and Ethes ???????????????? 43 Chapter 14 Aldehydes and Ketones ??????????????? 45 Chapter 15 Carboxylic Acids and Derivatives ????????? 48 Chapter 16 Amines, Other Nitrogen Compounds and Organic Sulfur
Compounds ???????????????????? 49 Chapter 17 Synthetic Polymers ????????????????? 53 Chapter 18 Carbohydrates ??????????????????? 56 Chapter 19 Proteins ?????????????????????? 59 Chapter 20 科技英语论文写作 ????????????????? 61
Chapter 1 Measurements
From simple chemicals the modern chemist can synthesize a drug with the ideal structural features to treat a particular disease or create a remarkable plastic with just the right properties to replace a worn
body part. Very rarely does a sudden, almost magical, discovery lead the way to this sort of success. In most cases careful, occasionally tedious, experimentation must come first. Performing experiments in chemistry and interpreting their results is what chemists do. It is with the devices used to produce measured quantities, the units in which they are expressed, and the techniques used to do calculations upon them that the study of chemistry begins.
Chemical experiments fall into two broad categories, qualitative experiments and quantitative
experiments. In qualitative experiments, the presence or absence of some physical quantity is noted. In quantitative experiments, the physical quantities are measured to see how much of it there is.
For example, in the experiments to test for glucose in urine, a Qualitative observation shows that the
sample contains glucose, whereas a Quantitative observation shows that the sample contains 10mg of glucose.
1.3 Units and the SI System
A unit describes a physical quantity that is being measured, e.g. 10 mg of glucose. A practical and useful set of units must be internationally accepted and unambiguously defined.
Three sets of units in use are:
a) English System: e.g., foot and pound, rarely used in scientific studies.
b) Metric system: e.g., meter and kilogram units, widely adopted.
c) International System of Units (SI System).
1.4 SI Units 1.4 SI Units
SI units were created in 1969, in order to clear up any possible confusion about which units should be included in the modern metric system. SI Units includes the SI base units, the SI derived units, and
the SI prefixes.
SI base units: There are seven SI base units for length, mass, time, amount of substance, temperature, electric current and luminous intensity, respectively. They are listed in the Table following.
Physical Quantity Name of Unit Abbreviation
Length meter m
Mass kilogram kg
Time second s
Amount of substance mole mol
Temperature kelvin K
Electric current ampere A
Luminous intensity candela cd
SI Prefixes: The nine widely used SI prefixes are listed in the following table.
Prefix Abbreviation Meaning
-12pico p 10 (one-trillionth)
-9nano n 10 (one-billionth)
-6micro 10 (one-millionth) ~
-3milli m 10 (one-thousandth)
-2centi c 10 (one-hundredth)
-1deci d 10 (one-tenth)
3kilo k 10 (one thousand times)
6mega M 10 (one million times)
SI Derived Units
SI derived units are in form of combinations of SI base units and, sometimes, SI Prefixes. For
example, Volume units are SI derived units.
3volume unit =(length unit)
31 cm = 1 ml (milliliter)
31 dm = 1 L (liter)
Containers that are used in chemical laboratories to measure volume include Beaker (量杯, 烧杯), Graduated cylinder (量筒), Burette (滴管), Syringe(注射器), Measuring pipet (吸量管 ), Transfer pipet(移液管),and Volumetric flask(容量瓶).
Another SI derived unit is for Density which is the combination of mass and volume units.
Density = mass/volume
3Water: 1.0 g/cm (1 g/ml)
3Gold: 19.3 g/cm (19.3 g/ml)
Chapter 2 Matter and Energy
Matter is anything that has mass and occupies space, e.g., a drink of water, a chunk of metal or even
a breath of air. Chemists study matter from one particular point of view, i.e., they explain the behavior of matter in terms of the invisible building blocks of which it is made. Atoms are the indivisible, discrete
particles of which all matter is composed. Molecules are collections of atoms which are held together by
links called chemical bonds.
2.2 Chemical Properties of Matter
A chemical property describes the ability of a substance to undergo a chemical change. A chemical
change occurs when the atoms of a substance rearrange by bond breaking and bond formation to
produce a new substance that is chemically different from the original ones. When such chemical
changes occur, a chemical reaction is said to have taken place. The original substances are called reactants and the new ones are called products.
Chemists describe chemical reactions by using an arrow pointing from the reactants to the products: Reactants ? products. Examples of chemical changes: burning, and corrosion.
2.3 Physical Properties of Matter
When a substance undergoes a physical change, no chemical bonds are formed or broken and no
chemical reaction takes place. The molecules and atoms of the original substance are the same before and after the physical change. Some important physical properties are color, density, boiling point and
2.4 States of Matter and Physical Changes
One physical property which is readily observed is the physical state, that is, whether something is a solid, a liquid, or a gas (at a given temperature and pressure).
Melting: the process that a solid is transformed into a liquid by applying heat to it.
Freezing: the reverse process of melting by cooling a liquid.
Vaporization: the process that a liquid is converted into a gas by heating.
2.5 Types of Matter
Elements: An element is a pure substance which consists of just one kind of atom. The 106 different elements on earth are listed in the Periodic Table of elements.
Compounds: A compound is a pure substance which contains just one kind of molecule.
Mixtures include more than one pure substance, which can be separated from each other without a chemical reaction.
2.6 Law of Conservation of Matter
Regardless of what chemical reaction takes place, careful weighings show that the mass of reactants is always exactly the same as the mass of products. Mass can not be created or destroyed, a principle
known as the law of conservation of matter.
Energy is defined as the ability to do work. Whirling tornadoes, rushing streams, and moving people are all sources of energy. The energy that involves objects in motion is called kinetic energy.
Stored energy is called potential energy, e.g., the water behind a dam. Chemical energy is the energy
change that accompanies chemical reactions.
Law of Conservation of Energy
The law of conservation of energy states energy is neither created nor destroyed, but may be changed in form.
2.8 Units of Energy
The SI energy unit is Joule, abbreviated J, a derived unit which is a combination of the kilogram,
meter, and second:
2-2 1 J = 1 kg m s
Chemists and biochemists sometimes substitute the non-SI energy unit calorie (cal):
1 cal = 4.184 J
a) Write an article (at least 200 words) with one of the following two topics:
Why is Chemistry Interesting?
Why is it necessary for you Chinese chemical students to learn Chemistry English?
Completing both will be highly appreciated!
Chapter 3 Atoms
In Greek atomos means “indivisible”.
Atomic theory: if the matter were divided a sufficient number of times, it could eventually be reduced to the indivisible, indestructible particles called atom.
The atomic theory was presented by the British chemist John Dalton (1766-1844) in the early 1800s.
It is one of the greatest advances in the history of chemistry. “Whether matter be atomic or not, this
much is certain, that granting it to be atomic, it would appear as it now does.”(by Micheal Faraday
(1794-1867) and J.B. Dumas(1800-1884)). The main points of the Atomic Theory include: 1) The
ultimate particles of elements are atoms; 2) Atoms are indestructible; 3) Elements consist of only one kind of atom; 4) Atoms of different elements differ in mass and in other properties; 5) Compounds consist of molecules (which Dalton called “compound atoms”), which form from simple and fixed combination of
different kinds of atoms.
Drawbacks of the Atomic Theory: Points 2 and 3 of Dalton‟s Atomic Theory do not agree with
modern experimental evidence because atoms can be broken down and atoms of one particular element can differ in mass.
3.2 Element Symbols
With the discovery of atoms came the chemical alphabet of element symbols. Dalton chose the circle as the symbol for oxygen and represented all other elements by variations of the circle. These early primitive symbols evolved into the modern system of using one or two letters of the English alphabet.
Modern system of element symbols: The first letter is always a capital and the second, if there is one, a lower case. The symbols are often formed from the first letter of the element name or from the first letter along with one other. e.g., B stands for the element boron, Ba for barium, Be for beryllium, and Bk
for berkelium.(锫, belongs to the Actinium(锕) series.)
Exceptions: For some of 106 elements it is not possible to guess the symbol by examining the English name. For instance, the symbol for the element iron is Fe (not I or Ir). Iron, along with copper, silver, gold, sodium, potassium, lead, tin, antimony, and tungsten have symbols that are derived from one or two letters of their Latin or German names.
The formulas used to represent compounds and elements inlcude element symbols and subscripts,e.g. HO represents a water molecule. 2
3.4 Subatomic particles
Particles smaller than even the smallest atoms are called subatomic particles. They are Electron
(1870s) , Proton (later 1800s) and Neutron (1930s).
3.5 Atomic mass unit (amu)
It is difficult to comprehend how incredibly small are the masses of subatomic particles. e.g. Proton
-24mass = 1.673 × 10 g
-24Neutron mass = 1.673 × 10 g
-28Electron mass = 9.11 × 10 g
Quoting the masses of these particles in grams is definitely awkward. A convenient unit to use is the
atomic mass unit.
-24 1 amu = 1.66057 × 10 g
3.6 Atomic Number Z
The identity of an element depends on the number of protons in the nuclei of its atoms. The number of protons in the nucleus of an atom is called the atomic number of the atom, labeled Z. All atoms of the same element must have the same number of protons. The number of positively charged protons and the
number of negatively charged electrons in an atom must be the same.
3.7 Isotopes and Mass Numbers 3.7 Isotopes and Mass Numbers
The sum of the number of protons and the number of neutrons in the nucleus of an atom is the mass
number. ( A = Z + N ). Atoms of the same element can have a different number of neutrons in
their nuclei. Isotopes are atoms of the same element which contain a different number of neutrons and
thus have different mass numbers.
Table Isotopes of Oxygen and Chlorine
Isotope p n e Natural Abundance, %
16O or O-16 8 8 8 99.76
17O or O-17 8 9 8 0.04
18O or O-18 8 10 8 0.20
35Cl or Cl-35 17 18 17 75.53
37Cl or Cl-37 17 20 17 24.47
3.8 Atomic Weight
Dalton recognized the hopelessness of ascertaining the absolute weights of atoms because atoms are
much too small to be weighted. It is possible to compare the weights of a large number of atoms of element A with that of the same number of atoms of element B. Atomic Weights for elements are
determined by comparing a very large number of the atoms of the element with the same number of
atoms of C-12. By definition the atomic weight of C-12 is exactly 12. For instance, the atomic weight of H is 1.008, meaning that H atoms are about one-twelfth as heavy as C-12 atoms.
Calculating the atomic weight
The atomic weight of an element is the weighted average of the atomic weights of all its natural isotopes and can be calculated if the atomic weights and relative abundances of the isotopes are given.
E.g., There are two naturally occurring chlorine isotopes, Cl-35 and Cl-37, with relative abundances of 75.5% and 24.5%, respectively.
Atomic Weight Cl = (0.755 × 35.0) + (0.245 × 37.0)
Atomic Weight Cl = 35.5
3.9 Formula Weight
The formula weight of an element or compound is calculated by adding the atomic weights all the
atoms in its formula.
e.g. Formula Weight of O = 2 × 16.0 = 32.0 2
Formula weight of HO = 2×1.0 + 1 × 16.0 = 18.0 2
3.10 Electrons in Atoms
It is the electrons that are responsible for the chemical properties of atoms. Electrons form the
bonds that connect atoms to one another to form molecules. The way in which the electrons are
distributed in an atoms is called the electronic structure of the atom. In an atom, the small, heavy positive nucleus is surrounded by circulating electrons.
n=4n=3n=23.11 Electronic Configurations n=1Each electron in an atom possesses a total energy (kinetic plus potential). The lowest-energy electrons are those closest to the nucleus of the atom and the most difficult to remove from the atom. Niels Bohr (1885-1962), a Danish physicist, first introduced the idea of electronic
n: rincial uantum number energy levels.
Bohr‟s Atomic Model was based on the Quantum Theory of Energy. The energy levels in atoms can
be pictured as orbits in which electrons travel at definite distances from the nucleus. These he called
“quantized energy levels”, also known as principal energy levels.
Schrödinger’s Atomic Theory
Bohr‟s theory laid the groundwork for modern atomic theory. In 1926, Erwin Schrödinger proposed
the modern picture of the atom, which is based upon a complicated mathematical approach and is used
today. In the Schrödinger atom, the principal energy level used by Bohr are further divided into sublevels, which are designated(指派；by a principal quantum number and a lowercase letter ( s, p, d and
f). The higher the energy level, the more sublevels there are. The electronic levels (1s, 2s,2p and so on)
are also called orbitals. zz
zz(and so on)
Atomic Orbitals Shapes of atomic orbitals:
s orbital is spherical;
p orbitals are dumbbell-shaped.