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General Chemistry I

Dr. PHAN TẠI HUÂN

Faculty of Food Science and TechnologyNong Lam University

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Recommended texts:– William L. Masterton and Cecile N. Hurley. 2009.

Chemistry: Principles and Reactions, 6th edition. Brooks/Cole Cengage Lerning.

– John A. Olmsted and Gregory M. Williams, 2005. Chemistry, 4th edition. John Wiley & Sons, Inc.

– Steven S. Zumdahl, 2005. Chemical Principles, 5th edition. Houghton Mifflin Company.

– Kenneth W. Whitten, Raymond E. Davis, Larry Peck, George G. Stanley, 2003. General Chemistry, 7th edition. Brooks Cole Publisher.

– Steven L. Hoenig, 2002. Basic Training in Chemistry. Kluwer Academic Publishers.

Course Introduction

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Attendance requirements• Class attendance: 45 hrs.

Assessment• Method Weighting

– Mid-term examination 30%– Final examination 70%

Course Introduction

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Why should we care about chemistry?– Chemistry is everywhere!

– Chemistry helps us to understand and be better informed about the world in which we live!

Course Introduction

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– Module 1: Atomic structure and Periodic Table.

– Module 2: Chemical bonds and molecular configurations.

– Module 3: The Three States of Matter.

– Module 4: Chemical Thermodynamics.

– Module 5: Chemical kinetics.

– Module 6: Solutions.

– Module 7: Oxidation-reduction reactions and transformation of chemical energy.

Course Introduction

Module 1: Atomic structure and Periodic Table

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Classification of Matter

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• Matter is composed of indivisible particles - atoms.

• An element is composed of only one kind of atom. These atoms in a particular element have the same properties such as mass, size, or even shape.

• A compound is composed of two or more elements combined in fixed ratios or proportions.

• In a chemical reaction, the atoms in the reactants recombine, resulting in products which represent the combination of atoms present in the reactants. In the process, atoms are neither created, nor destroyed. So a chemical reaction is essentially a rearrangement of atoms.

Dalton's Atomic Theory (1803)

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• The Law of Conservation of Mass states that mass is neither created nor destroyed in a chemical reaction or physical change.

• The Law of Definite Proportions states that every chemical compound is made up of elements in a definite ratio by mass.

Ramifications of Dalton's Theory

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Ramifications of Dalton's Theory

Exercise 1: Calculate the mass of sodium chloride formed by the complete reaction of 10.0 g of sodium with 15.4 g of chlorine. What law allows this calculation?

• Ans:

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• Exercise 2: Calculate the mass of oxygen that will combine with 2.00 g of magnesium if 0.660 g of oxygen reacts with 1.00 g of magnesium. What law allows this calculation?

• Ans:

Ramifications of Dalton's Theory

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Schematic drawing of a gas discharge tube in operation. When a high voltage is applied to the two perforated plates, an electrical discharge occurs between them. The positively charged and negatively charged particles that form in the gas move to collectors at the ends of the tube.

The discovery of the electron

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J. J. Thomson experiment (1897)

• Schematic drawing of a cathode-ray tube. A beam of electrons is deflected by a pair of charged plates (bent line), but magnetic force can be adjusted to exactly counterbalance the effect of the electrical force (straight line).

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Schematic view of Millikan's oil drop experiment. An atomizer generated a fine mist of oil droplets. Bombarding the droplets with X rays gave some of them extra negative charge. In the presence of sufficient electrical force, these negatively charged droplets could be suspended in space.

Millikan's oil drop experiment (1909)

• The measurements gave several different values, but the charge was always equal to n(-1.6 x 10-19C).

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Thomson´s plum pudding model

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Rutherford's scattering experiment

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The general structrure of the atom

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• What is the net charge on an atom that contains 8 protons, 8 neutrons, and 10 electrons?

• Ans:

Exercise

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• The atomic number (Z) denotes the number of protons in an atom's nucleus.

• The mass number (A) denotes the total number of protons and neutrons. Protons and neutrons are often called nucleons.

• Some atoms have the same atomic number, but different mass numbers. This means different number of neutrons. Such atoms are called isotopes.

Constituents of the atom

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Isotopes

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Exercise

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Periodic chart of the elements

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• Atomic mass unit (amu) is defined as exactly 1/12 the mass of a 12C atom. The mass of the 12C atom is taken to be exactly 12 amu.

Atomic mass unit (amu)

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• The atomic mass of an element is the weighted average of the masses of the individual isotopes of the element.

• Example: Naturally occurring copper consists of 69.17% 63Cu and 30.83% 65Cu. The mass of 63Cu is 62.939 598 amu, and the mass of 65Cu is 64.927 793 amu. What is the atomic mass of copper?

Atomic mass

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• Naturally occurring carbon consists of two isotopes, 12C and 13C. What are the percentage abundances of the two isotopes in a sample of carbon whose atomic mass is 12.01112?

• Ans:

Exercise

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• The quantity of a given substance that contains as many units or molecules as the number of atoms in 12 grams of carbon-12 is called a mole.

• The numerical value of one mole is 6.023 x 1023 and is referred to as Avogadro´s number.

• One mole of hydrogen atoms contains Avogadro number of hydrogen atoms.

The concept of mole

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Atomic structure

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In 1864, James Clerk Maxwell developed a mathematical theory to describe radiation as wave-like, or oscillating, electric and magnetic fields in space. The electric and magnetic fields are perpendicular to each other.

Electromagnetic radiation

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• Visible light, infrared radiation, microwaves, radio waves, ultraviolet, x-rays and gamma rays are all types of electromagnetic radiation.

Electromagnetic spectrum

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• The distance between two waves, usually measured from the peak of the waves, is the wavelength, given the symbol lambda, λ.

• The frequency is a statement of the number of waves passing a point in space per second; it is given the symbol nu, ν.

(The hertz is commonly used as the unit for frequency; 1 Hz = 1 s−1)

Wave character of light

• The product of the wavelength and the frequency is equal to the velocity of light, usually designated by c:

c = λν(The value of c can be rounded to c=2.998 × 108 m/s for most calculations.)

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• The energy of light is emitted, absorbed, or converted to other forms of energy in individual units referred to as quanta (singular: quantum).

• The unit of light energy is often referred to as the particle of light called the photon.

• The energy of a photon is proportional to the frequency:ε = hν = (6.626 × 10−34 J · s)ν

Planck’s constant, h, is the universal proportionality constant.

Particle character of light

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• When a metal bar is heated, it emits electromagnetic radiation observable as the red to white glow of the metal. What is the energy of one photon of red light with a wavelength of 700 nm? What is the energy of a mole of photons with this wavelength? Calculate the energy of one photon and a mole of photons of blue light with a wavelength of 400 nm.

Exercise

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Exercise

• What is the wavelength in meters of the radiation from (a) a low-range TV station broadcasting at a frequency of 55 MHz, (b) anAM radio station at 610 kHz, and (c) a microwave oven operating at 14.6 GHz?

• Ans:

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Interaction of light with matter

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• In the late 19th century, Johann Balmer (1825–1898) and Johannes Rydberg (1854–1919) showed that the wavelengths of the various lines in the hydrogen spectrum can be related by a mathematical equation:

• Here R is 1.097 x 107 m-1 and is known as the Rydberg constant. The n’s are positive integers, and n1 is smaller than n2

• However this is an empirical equation.

Atomic spectra

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1) In each hydrogen atom, the electron revolves around the nucleus in one of the several stable orbits.

2) Each orbit has a definite radius and thus has a definite energy associated with it.

3) An electron in an orbit closest to the nucleus has the lowestenergy, and if the electron is in the lowest orbit the atom is said to be in its ground state.

4) The electron in an atom may absorb discrete amounts of energy and move to another orbit with higher energy, and this state is called the excited state.

5) An electron in an excited atom can go back to a lower energy level and this process will result in the release of excess energy as light.

6) The amount of energy released or absorbed.

Bohr's Model of Hydrogen Atom (1913)

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The radius is given byr= n2a0

a0 =5.292 x 10-11 m = 0.5292 Å (Bohr radius)

The potential energy is given by:

where h Planck’s constant, m the mass of the electron

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• Each line in the emission spectrum represents the difference in energies between two allowed energy levels for the electron.

• Comparing this to the Balmer-Rydberg equation

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Bohr's Model of Hydrogen Atom

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• De Broglie (1925) proposed that not only does light have the dual properties of waves and particles, but also particles of matter have properties of waves. The wavelength of those particle waves is given by

λ = h/ mv

where m and v are the mass and velocity of the particle. Planck’s constant, h, is so small that the wavelengths are in an observable range only for particles of atomic or subatomic mass.

The wave nature of the electron

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• The Heisenberg Uncertainty Principle, stated in 1927 by Werner Heisenberg (1901–1976):

It is impossible to determine accurately both the momentum and the position of an electron (or any other very small particle) simultaneously.

Quantum mechanical picture of the atom

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• Atoms and molecules can exist only in certain energy states. In each energy state, the atom or molecule has a definite energy. When an atom or molecule changes its energy state, it must emit or absorb just enough energy to bring it to the new energy state (the quantum condition).

• When atoms or molecules emit or absorb radiation (light), they change their energies.The energy change in the atom or molecule is related to the frequency or wavelength of the light emitted or absorbed by the equations:

• The allowed energy states of atoms and molecules can be described by sets of numbers called quantum numbers.

Basic ideas of quantum mechanics

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• The mathematical approach of quantum mechanics involves treating the electron in an atom as a standing wave

Basic ideas of quantum mechanics

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• A solution to the Schrödinger equation for an electron must satisfy three quantum conditions corresponding to the three dimensions of space.

• Each quantum condition introduces an integer, called a quantum number, into the solution.

• A separate solution, describing a probability distribution of finding the electron at various locations, exists for each allowed set of three quantum numbers.

• Such a solution is called an orbital.

Orbitals

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• Principal quantum number (n). The principal quantum number denotes the energy level of electrons. The larger the principal quantum number is, the larger the energy.

• The orbital size depends on n. This means that the larger the n value, the larger the orbital. Orbitals with the same n belong to the same shell.

Quantum number

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• Angular momentum quantum number (l). Angular momentum quantum number denotes the shape of the orbital. The values range from 0 to n – 1.

• The angular momentum quantum numbers correspond to different subshells.

Quantum number

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• Magnetic quantum number (ml). Magnetic quantum numbers define the different spatial orientations of the orbitals. The values range from –l to +l.

• There are three p orbitals corresponding to ml = 1, 0, and −1. However, it is usually more convenient in chemistry to use a new set of three orbitals oriented along the x, y, and z axes to display the shapes and directions of these orbitals.

• Further, there are 5 d orbitals and 7 f orbitals having different shapes and orientations in space.

Quantum number

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Orbitals

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Orbitals

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• Spin quantum number (mS) Spin quantum number has to do with the spin orientations of an electron. The two possible spins are denoted by the spin quantum numbers + ½ and -1/2.

Quantum number

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• The values of n, l, and ml describe a particular atomic orbital. Each atomic orbital can accommodate no more than two electrons, one with ms=+1/2 and another with ms=-1/2 .

• A set of quantum numbers: (n, l, ml, ms).

Quantum number

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Permissible Values of the Quantum Numbers Through n=4

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• No two electrons in an atom may have identical sets of four quantum numbers (Pauli exclusion principle).

• Orbitals are filled in the order of increasing energy (Klechkowski´s rule).

• Electrons occupy all the orbitals of a given subshell singly before pairing begins. These unpaired electrons have parallel spins (Hund´s rule).

“Building” of Electron configurations by the Aufbau Principle

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Two general rules help us to predict electron configurations.

• Electrons are assigned to orbitals in order of increasing value of (n + l ).

• For subshells with the same value of (n+l ), electrons are assigned first to the subshell with lower n.

1s < 2s < 2p <3s < 3p<4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d...

Electron configurations

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Exercise

• Arrange the electrons in the following list in order of increasing energy, lowest first:

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Exercise

• How many electrons are permitted in each of the following subshells? (a) 2s (b) 6p, and (c) 4d.

• Ans:

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Exercise

• How many electrons are permitted in each of the following subshells?

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• Write the electronic configuration of sulfur and also show the filling of electrons with orbital notation.

• Ans:

Exercise

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• Write detailed electronic configurations for (a) N (Z=7), (b) P (Z=15), (c) As (Z=33), and (d) Sb (Z=51).

• What makes their chemical properties similar?

• Ans:

Exercise

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The periodic table

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• The properties of elements are periodic functions of their atomic numbers

• The vertical columns of elements represented in the periodic table are called groups, and the horizontal rows are called periods.

• There are seven periods in the periodic table. The groups are usually designated by roman numerals followed by the letter A or B as shown in the periodic table.

The periodic table

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• The groups IA through VIIA are called the representative elements. These elements have either s or p orbital valence electrons. The last group in the periodic table is the noble gas group otherwise known as the zero group.

• The groups ranging from IB through VIIIB are called transition metals, and finally the metals from lanthanum through hafnium and metals from actinium onward are called the inner transition metals.

The periodic table

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Period 1

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Period 2

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Period 3

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Period 4

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As principal quantum number n increases, atomic orbitals become larger and less stable. As atomic number Z increases, any given atomic orbital becomes smaller and more stable.

Periodicity of Atomic Properties

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Exercise: Arrange the following elements in terms of increasing atomic radius: Mg, Cl, K and Cs.

Atomic size

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Atomic size

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• Ionization energy (IE) is the minimum amount of energy required to remove an electron from an atom.

Ionization energy

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Ionization energy

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• First IE < Second IE < Third IE < Fourth IE < ...

Ionization energy

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• Electron affinity is the energy change associated with the addition of an electron to a gaseous atom.

Electron affinity

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Electron affinity

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Electron affinity

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Electronegativity

• The relative tendency of an atom to attract the bonding electrons to itself is called electronegativity. The popularly used electronegativity scale is based on a system called Pauling's scale, according to which fluorine (the most electronegative element) has an electronegativity value of 4.0. Nonmetals are the most electronegative elements.

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Electronegativity Values of the Elements

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After you have studied this module, you should be able to

• Describe the evidence for the existence and properties of electrons, protons, and neutrons.

• Predict the arrangements of the particles in atoms.

• Describe isotopes and their composition.

• Calculate atomic weights from isotopic abundance.

• Descriptions of waves play an important role in our theories of light and of atomic structure.

Summary

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• Describe the four quantum numbers, and give possible combinations of their values for specific atomic orbitals.

• Describe the shapes of orbitals and recall the usual order of their relative energies.

• Write the electron configurations of atoms.

• Relate the electron configuration of an atom to its position in the periodic table.

• Periodicity of atomic properties.

Summary