Course code : CHEM 11122 Course title : General Chemistry and Basic Analytical Chemistry

Course content: Structure and bonding (10 h) Aqueous Solution Chemistry (20 h)

Observations in the macroscopic world

Matter is composed of atoms. Understanding the structure of atoms is critical to understand the properties of matter

History of atomic theory

A revision of basic concepts

A Brief History of Atomic Theory

This section will focus on scientists who have had an impact on the study of the atom.

Democritus Democritus proposed that matter cannot be broken down indefinitely.

470-380 B.C.

At some point you end up with a piece that can’t be divided. That smallest piece he called an atom, from the Greek word atomos, which means “indivisible”.

Democritus’ Model

“Atomos” ATOMOS was the word Democritus used the point, or stage cannot be broken down any ATOMOS waswhere thematter word Democritus used further. ATOMOS literally means “indivisible”

the point, or stage where matter cannot be broken down any further. ATOMOS literally means “indivisible”

The Birth of Modern Atomic Theory

John Dalton • John Dalton was a British chemist. • He was the first modern scientist to propose the existence of atoms. • He described an atom as an invisible indestructible, solid sphere, like a billiard ball.

1766 - 1844

Dalton’s Model

The “Indivisible Sphere”

John Dalton • In 1803, proposed an Atomic Theory which states: -Elements are made out of minute, indivisible particles –Atoms -Atoms of the same element are exactly alike, and atoms of different elements are different - Atoms can not be created or destroyed

Image taken from: chemistry.about.com/.../JohnDalton.htm

-Compounds are formed union of two or more atoms of different elements in a simple whole number ratios.

Subatomic Particles

J.J. Thomson (1856 – 1940)

• J.J. Thomson was a British physicist • Proved that an atom can be divided into smaller parts

Image taken from: www.wired.com/.../news/2008/04/d ayintech_0430

• While experimenting with cathode-ray tubes, discovered corpuscles, which were later called electrons

•Stated that the atom is neutral •In 1897, proposed the Plum Pudding Model which states that atoms mostly consist of positively charged material with negatively charged particles (electrons) located throughout the positive material

Thomson “plum pudding” model

+

atoms mostly consist of positively charged material with negatively charged particles (electrons) located throughout the positive material

A = alpha B = gamma C = beta

J.J. Thomson, measured charge/mass of e(1906 Nobel Prize in Physics)

Rutherford’s experiment (1871-1937)

Experimented with a radiation source that sent out alpha particles through a thin piece of gold foil to a detector screen that glowed when it was hit.

•Suggested the following characteristics of the atom: oIt consists of a small core, or nucleus, that contains most of the mass of the atom oThis nucleus is made up of particles called protons, which have a positive charge oThe protons are surrounded by negatively charged electrons, but most of the atom is actually empty space

Ernest Rutherford experiments proved that atoms are mostly empty space. Discovered the nucleus, which contains positively charged particles .

1932 A.D. English scientist

James Chadwick Using alpha particles discovered a neutral atomic particle with a mass close to a proton. What he discovered was the neutron. -Discovered the third subatomic particle of the atom: the neutron 450 BC 400 BC 350 BC

500-1600 AD

1650 AD Late 1700’s 1808AD 1831AD 1879AD1897AD1898AD

1911AD

1922AD

1932AD

1922 A.D. Danish scientist

Niels Bohr Came up with experimental evidence proving that electrons exist in energy levels (shells) orbiting around a positively charged nucleus.

450 BC 400 BC 350 BC

500-1600 AD

1650 AD Late 1700’s 1808AD 1831AD 1879AD1897AD1898AD

1911AD

1922AD

1932AD

The arrangement of electrons in an atom is called its electronic structure. Much of our present understanding of the electronic structure of atoms has come from analysis of the light emitted or absorbed by substances. -To understand the basis for current model of electronic structure we must first learn more about light

Light and atomic spectra

Light -The light we see with our eyes, visible light, is one type of electromagnetic radiation. -There are many types of electromagnetic radiation in addition to visible light. -radio waves that carry music to our radios, -infrared radiation (heat) from a glowing fireplace, -X-rays -All types of electromagnetic radiation move through a vacuum at , 3.00 x 108 m/s the speed of light.

-All have wave-like characteristics similar to those of waves that move through water.

-The distance between two adjacent peaks (or between two adjacent troughs) is called the wavelength - The number of complete wavelengths, or cycles, that pass a given point each second is the frequency of the wave

There is an inverse relationship between the frequency and wavelength of electromagnetic radiation

Practice Exercise 1

Electromagnetic Spectrum -shows

the various types of electromagnetic radiation arranged in order of increasing wavelength

the wavelengths of radio waves can be longer than a football field

The wavelengths of gamma rays are comparable to the diameters of atomic nuclei

Wave-like Behaviors of Light How Do We Know that Light is a Wave? Light reflects in the same manner that any wave would reflect. Light refracts in the same manner that any wave would refract. Light diffracts in the same manner that any wave would diffract.

Since light behaves like a wave, one would have good reason to believe that it might be a wave

-Although the wave model of light explains many aspects of its behavior, this model cannot explain several phenomena

(1) the emission of electrons from metal surfaces on which light shines (the photoelectric effect),

(2) the emission of light from electronically excited gas atoms (emission spectra)

In 1900 a German physicist named Max Planck (1858–1947)

energy can be either released or absorbed by atoms only in discrete amounts called “quanta”

quantum = (meaning “fixed amount”) to the smallest quantity of energy that can be emitted or absorbed as electromagnetic radiation.

He proposed that the energy, E, of a single quantum equals a constant times the frequency of the radiation:

-constant h is called Planck’s constant

and has a value of = 6.626 x 10-34joule second (J s).

-Because the energy can be released only in specific amounts, we say that the allowed energies are quantized -Planck’s revolutionary proposal that energy is quantized was proved correct, and he was awarded the 1918 Nobel Prize in Physics for his work on the quantum theory

In 1905, Albert Einstein (1879–1955) used Planck’s theory to explain the photoelectric effect

The Photoelectric Effect

The photoelectric effect is the observation that many metals emit electrons when light shines upon them.

A minimum frequency of light, different for different metals, is required for the emission of electrons

-the radiant energy striking the metal surface behaves like a stream of tiny energy packets. Each packet, which is like a “particle” of energy, is called a photon Light has particle like properties

Under the right conditions, photons striking a metal surface can transfer their energy to electrons in the metal. A certain amount of energy—called the work function —is required for the electrons to overcome the attractive forces holding them in the metal. If the photons striking the metal have less energy than the work function, the electrons do not acquire sufficient energy to escape from the metal, even if the light beam is intense.

If the photons have energy greater than the work function of the particular metal the excess appears as the kinetic energy of the emitted electrons

Einstein won the Nobel Prize in Physics in 1921 for his explanation of the photoelectric effect.

Einstein’s theory of light as a stream of photons rather than a wave

Light has both wave and particle like properties

Practice Exercise 2

Line Spectra In 1913, the Danish physicist Niels Bohr offered a theoretical explanation of line spectra, another

phenomenon that had puzzled scientists during the nineteenth century.

spectrum A spectrum is produced when radiation from sources( light bulbs and stars) is separated into its component wavelengths

A continuous visible spectrum is produced when a narrow beam of white light is passed through a prism. The white light could be sunlight

Not all radiation sources produce a continuous spectrum.

Atomic emission of hydrogen and neon Different gases emit light of different characteristic colors when an electric current is passed through them.

When light coming from such tubes is passed through a prism, only a few wavelengths are present in the resultant spectra

Line Emission Spectrum of Hydrogen Atoms

four narrow bands of bright light are observed against a black background

Line spectrum of neon

Each colored line in such spectra represents light of one wavelength. Every element has a unique emission spectrum A spectrum containing radiation of only specific wavelengths is called a line spectrum.

Atomic Spectrum of Hydrogen

• In 1885 Johann Balmer showed that the wave length of any line in the visible spectrum of atomic hydrogen, could be given by the simple formula;

Later, additional lines were found in the ultraviolet and infrared regions of hydrogen’s line spectrum.

Soon Balmer’s equation was extended to a more general one, called the Rydberg equation, which allows us to calculate the wavelengths of all the spectral lines of hydrogen

How could the remarkable simplicity of this equation be explained?

Bohr theory To explain the line spectrum of hydrogen, Bohr assumed that electrons in hydrogen atoms move in circular orbits around the nucleus like planets around the sun

Several Problems arise with this concept 1. Electrons are expected to slow down gradually. 2. Why does electrons move in an orbit around nucleus? 3. Since nucleus and electrons have opposite charges, they should attract each other and collide.

To explain these problems Bohr postulates 1. Electrons do not radiate energy if stay in one orbit, ∴ do not slow down. 2. When electrons move from one orbit to another they radiate or absorb energy. 3. For an electron to remain in orbit the electrostatic attraction between the nucleus and electron must be equal to the centrifugal force.

The Bohr Model of the Atom

Electronic transition and origin of the spectral lines of hydrogen atom

Bohr assumed that the discrete lines seen in the spectrum of the hydrogen atom were due to transitions of an electron from one allowed orbit/energy level to another.

What is most significant about Bohr’s model is that it introduces two important ideas;

1. Electrons exist only in certain discrete energy levels, which are described by quantum numbers. 2. Energy is involved in the transition of an

electron from one level to another.

Limitations of Bohr's model of the atom -The Bohr model could only successfully explain the hydrogen spectrum. -It could NOT accurately calculate the spectral lines of larger atoms. -The model only worked for hydrogenlike atoms That is, if the atom had only one electron.

Depending on the experimental circumstances, radiation appears to have either a wave-like or a particle-like (photon) character.

Wave-particle duality of matter • In1924, Prince Louis Victor de Broglie(18921987)- French physicist- proposed that all particles of matter (from single atoms to large objects) moving at some velocity would have the properties of a wave.

De Broglie used the term matter waves to describe the wave characteristics of material particles.

The Heisenberg Uncertainty Principle • In 1927,Werner Heisenberg(1901-1976) German physicist- no experimental method

can be devised that will measure simultaneously the precise position as well as the momentum of an object.

The Heisenberg Uncertainty Principle • It is impossible to know simultaneously both the position and the momentum of an object/electron exactly.

• ∴ The probability of finding an electron in a particular position or in a particular volume of space has to be considered.

Practice Exercise 4

Thus, we have essentially no idea where the electron is located in the atom.

• ∴ The probability of finding an electron in a particular position or in a particular volume of space has to be considered.

QUANTUM MECHANICS AND ATOMIC ORBITALS Quantum mechanics was used in 1926 by Erwin Schrödinger(1887-1961) – Austrian physicist – to describe electrons in atoms. -proposed an equation, now known as Schrödinger’s wave equation Solving Schrödinger’s equation for the hydrogen atom leads to a series of mathematical functions called wave functions.

-These wave functions are usually represented by the symbol (lowercase Greek letter psi). -the square of the wave function, at a given point in space represents the probability that the electron will be found at that location. For this reason, is called either the probability density or the electron density.

Where in the figure is the region of highest electron density?

-Electron-density distributionThis rendering represents the probability of finding the electron in a hydrogen atom in its ground state

Quantum numbers for electrons The Bohr model was a one-dimensional model that used one quantum, number to describe the distribution of electrons in the atom/ orbit. The quantum mechanical model uses three quantum numbers, n, l, and , ml which result naturally from the mathematics used, to describe the distribution of electrons in the atom/orbitals

Quantum numbers for electrons 1. Principal Quantum Number (n) n = integral values of 1, 2, 3, 4, ……. K L M N …….. Determines the main energy level (shell) the electron is in. As n increases, the orbital becomes larger, and the electron spends more time farther from the nucleus. An increase in n also means that the electron has a higher energy and is therefore less tightly bound to the nucleus.

2. Azimuthal (angular momentum) Quantum Number (l) l = 0, 1, 2, 3,…….(n-1)

-Describes the sub-shell that the electron occupies. -This quantum number defines the shape of the orbital.

3. Magnetic Quantum Number (ml) can have integral values between -l and l, including zero. ml = +l…..0…..-l This quantum number describes the orientation of the particular orbital that the electron occupies in space

•For s sublevel, l = 0

m=0

1 s type

•For p sublevel, l = 1 • •

3 p types

m = +1, 0, -1

•For d sublevel, l = 2

m = +2,+1, 0, -1,-2

• 5 d types

•For f sublevel, l = 3

m = +3,+2,+1, 0, -1,-2,-3

7 f types

Practice problem 05

4. Spin Magnetic Quantum Number (ms) ms = +1/2 or –1/2

-The electron behaves as if it were spinning about an axis, thereby generating a magnetic field whose direction depends on the direction of spin. -The two directions for the magnetic field correspond to the two possible values for the spin quantum number

Zeeman effect •In 1896 Pieter Zeeman (1865-1943) - Dutch physicist - discovered that the spectral lines of a light source subjected to a strong magnetic field were split into several components. •(This phenomenon, known as the Zeeman effect.

A photo Zeeman took of the Zeeman effect

Comparing probability density and radial probability function .

Radial probability distributions for the 1s, orbital of hydrogen

Comparison of the 1s, 2s, and 3s orbitals

The Pauli exclusion principle In 1925 the Austrian-born physicist Wolfgang Pauli (1900–1958) discovered the principle that governs the arrangements of electrons in manyelectron atoms. The Pauli exclusion principle states that no

two electrons in an atom can have the same set of four quantum numbers n, l, ml and ms. -an orbital can hold a maximum of two electrons and they must have opposite spins.

PRACTICE EXERCISE 06

ELECTRON CONFIGURATIONS The way electrons are distributed among the various orbitals of an atom is called the electron configuration of the atom.

the orbitals are filled in order of increasing energy, with no more than two electrons per orbital.

A half arrow pointing up represents an electron with a positive spin magnetic quantum number and a half arrow pointing down represents an electron with a negative spin magnetic quantum number .

Electrons having opposite spins are said to be paired when they are in the same orbital. An unpaired electron is one not accompanied by

a partner of opposite spin.

In the lithium atom the two electrons in the 1s orbital are paired and the electron in the 2s orbital is unpaired.

Order of Filling Orbitals Hund’s rule •Within a sublevel,each orbital is occupied by

one electron before any orbital has two(pairing). •The first electrons to occupy orbitals with in a

sublevel must have parallel spins. Hund’s rule, which states that for

degenerate orbitals, the lowest energy is attained when the number of electrons having the same spin is maximized.

Aufbau principle

General energy ordering of orbitals for a many-electron atom.

•The filling begins with the sublevel lowest in energy and continues upwards according to the “Aufbau principle”

PRACTICE EXERCISE 07

Condensed Electron Configurations -In writing the condensed electron

configuration of an element, the electron configuration of the nearest noble-gas element of lower atomic number is represented by its chemical symbol in brackets.

Anomalous Electron Configurations The electron configurations of certain elements appear to violate the rules we have just discussed This anomalous behavior is largely a consequence of the closeness of the 3d and 4s orbital energies. It frequently occurs when there are enough electrons to form precisely half-filled sets of degenerate orbitals (as in chromium) or a completely filled d subshell (as in copper).

ELECTRON CONFIGURATIONS AND THE PERIODIC TABLE

Dmitri Mendeleev In 1869 he published a table of the elements organized by increasing atomic mass

1834 - 1907

The Modern Periodic Table Element Organization The Periodic Law: Elements are arranged by atomic #,

Periodic table of the elements (columns) Groups

Periods (rows)

Elements are presented in the periodic table by increasing values of their atomic numbers, the number of protons in their atomic nuclei

Symbols in Table

periodic table blocks -named according to the subshell in which the "last" electron resides.

Chemical compounds A compound is a substance containing more than one element

CHEMICAL BONDS Two of the most common substances on our dining table are salt and granulated sugar

NaCl

C12H22O1 1

The properties of substances are determined in large part by the chemical bonds that hold their atoms together

CHEMICAL BONDS
Atoms or ions are held together in molecules or compounds by chemical bonds.

TYPES OF CHEMICAL BONDS

• Ionic bonds • Covalent bonds • Metallic bonds • Coordinate bond

Ionic Bond • Refers to electrostatic forces that exist between ions of opposite charges. • Between atoms of metals and nonmetals with very different electronegativity • Conductors and have high melting point. • Examples; NaCl

-Ionic bond – electron from Na is transferred to Cl. -The Na becomes (Na+) and the Cl becomes (Cl-), charged particles or ions.

Atoms often gain, lose, or share electrons to achieve the same number of electrons as the noble gas closest to them in the periodic table.

The noble gases have very stable electron arrangements

Octet rule:

Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons.

Metallic Bond

Metallic Bond • Metallic bonding constitutes the electrostatic attractive forces between the delocalized electrons, called conduction electrons, gathered in an electron cloud and the positively charged metal ions • Good conductors at all states, • Examples; Na, Fe, Al, Au, Co

Covalent Bond • Between nonmetallic elements of similar electronegativity. • Formed by sharing electron pairs • Stable non-ionizing particles, they are not conductors at any state • Examples; O2, CO2, C2H6, H2O,

Lewis Structures -we

usually show each shared electron pair as a line and any unshared electrons as dots.

- the Lewis structures for H2 and Cl2 are

MOLECULAR GEOMETRY AND BONDING THEORIES

Valence Shell Electron Pair Repulsion (VSEPR) Theory 1. Shapes of molecules are determined by the repulsions between electron pairs in the valence shell 2. A lone pair takes up more space around the central atom than a bond pair. Hence repulsion between lone pair – lone pair > bond pair-lone pair > bond pair- bond pair

The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.

Trigonal planar

Bent

0

1

2

Bent

Valence Bond Theory/VB theory -Covalent bonding occurs when a valence atomic orbital of one atom overlaps, with a valence atomic orbital of another atom. -The overlap of orbitals allows two electrons of opposite spin to share the space between the nuclei, forming a covalent bond.

Covalent bonds in H2, HCl, and Cl2 result from overlap of atomic orbitals

Potential Energy Curve

-shows how the potential energy of a system consisting of two H atoms changes as the atoms come together to form an H2 molecule

-When the atoms are infinitely far apart, they do not “feel” each other and so the energy approaches zero. -As the distance between the atoms decreases, the overlap between their 1s orbitals increases (the potential energy of the system decreases. That is, the strength of the bond increases; attractive forces dominate

-shows that as the atoms come closer together than 0.74 Å, the energy increases sharply (is due mainly to the electrostatic repulsion between the nuclei)

-The internuclear distance, or bond length, is the distance that corresponds to the minimum of the potential-energy curve (This is moe correctly known as the equilibrium bond length). -The potential energy at this minimum corresponds to the bond strength.

Sigma bonds •End to end or head on overlap of orbitals result in bonds. •The electron density is concentrated in between the two atoms. -the line joining the two nuclei passes through the middle of the overlap region. These bonds are called sigma bonds.

Pi (П) bonds • Multiple bonds form by the sideway overlap of orbitals, known as p orbitals. • A bond is one in which the overlap regions lie above and below the internuclear axis.

Hybridization and hybrid orbitals The process of mixing atomic orbitals is a mathematical operation called hybridization.

Formation of two equivalent Be-F bonds in BeF2.

How many atomic orbitals contribute to form the three sp2 hybrid orbitals?

Formation of sp2 hybrid

Formation of sp3 hybrid orbitals

Molecular orbital theory Molecular orbital theory describes the electrons in molecules by using specific wave functions called molecular orbitals (MO). The Hydrogen Molecule

The two molecular orbitals of H2, one bonding MO and one an antibonding MO.

-In the bonding MO electron density is concentrated in the region between the two nuclei. -By contrast, the antibonding MO has very little electron density between the nuclei. -The relative energies of two 1s atomic

orbitals and the molecular orbitals formed from them are represented by an energy-level diagram (also called a molecular orbital diagram

Energy-level diagrams and electron configurations for H2