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Presentation on theme: "GENERAL CHEMISTRY University POLITEHNICA Timisoara"— Presentation transcript:

Faculty of Industrial Chemistry and Environmental Engineering GENERAL CHEMISTRY Andrea Kellenberger

2 GENERAL CHEMISTRY Lecture: 2h / week Laboratory: 2h / week
Evaluation form: Examination Nr. of credits: 5

3 INTRODUCTION Chemistry is the science of substances.
Chemistry studies the structure, properties and transformation of substances. Substances are charaterized by 2 essential properties: homogeneity and constant composition.

4 INTRODUCTION Homogeneity = is the property of substances to show same characteristics in all the volume. Constant composition = in any given portion of a substance there are the same particles which interact in the same way. Substances can not be separated into other substances by physical methods.

5 INTRODUCTION Exemple of substances: water, sugar, oxygen, hydrogen, sodium chloride, copper, hydrochloric acid, sodiu hydroxide. Air = substance? Air is not a substance. By distillation of liquefied air one obtains O2, N2 and other gases.

6 INTRODUCTION Gasoline = substance?
Gasoline is not a substance, because by distillation it can be separated into the component hydrocarbons. NaCl solution = substance? Sodium chloride solution is not a substance, because by boiling the water is removed and the NaCl crystals remains.

7 INTRODUCTION In nature, substances are found only rarely in pure form.
Materials are made ​​of substances or mixtures of substances. Materials can be homogeneous sau heterogeneous. Heterogeneous materials are composed of so-called "phases" which are homogeneous portions separated from other phases by areas where the properties vary sharply.

8 INTRODUCTION Exemple of materials:
metals and alloys. Metals are homogeneous materials containing a single metallic element. Alloys are homogeneous or heterogeneous mixtures of two or more metals. Alloys: Brass = copper and zinc alloy Bronze = copper-tin alloy

9 INTRODUCTION Exemple of materials:
Wood = homogeneous material, composed of cellulose, lignin, resins and other substances. Glass = usually a homogeneous amorphous material (component particles are not included in a crystal lattice) Granite = homogeneous material composed of three different minerals: quartz, feldspar and mica.

10 INTRODUCTION Solutions = homogeneous mixture composed of two or more substances. Solutions are obtained by dissolving a substance called solute into another substance called solvent. Exemples liquid solutions solid solutions gaseous solutions

11 INTRODUCTION Liquid solutions: gas dissolved in liquid (CO2 in water)
liquid dissolved in liquid (ethanol in water) solid dissolved in liquid (copper sulfate in water, naphtalene in benzene) Solid solutions = homogeneous alloys Gaseous solutions = homogeneous mixtures of gases, i.e. air. <a href="">Bare de aur</a> de Petr Kratochvil

12 INTRODUCTION Each substance is distinguished from the other substances by its properties. A substance is characterized by its physical properties that can be expressed by numerical values. For pure substances these properties are called physical constants. melting point, boiling point density refractive index dielectric constant

13 International System of Units (S.I.). S.I. has 7 fundamental units
INTRODUCTION International System of Units (S.I.). S.I. has 7 fundamental units Quantity Unit Symbol Length meter m Mass kilogram kg Time second s Thermodynamic temperature Kelvin K Electric current Ampere A Luminous intensity candela cd Amount of substance mole mol

14 Derived units Quantity Equation Name of the unit Symbol Area m2 Volume
S = l2 l – length square meter m2 Volume V = l3 cubic meter m3 Speed v = l / t t – time meter per second m s-1 Aceleration a = v / t meter per second squared m s-2 Molar concentration cM = n / V n- amount of substance V-volume mole per cubic meter mol m-3 Force F = m·a m – mass a – acceleration kilogram· meter per second squared (Newton) kg· m s-2( N ) Pressure P = F / S S – area kilogram· meter per square second per square meter kg· m s-2· m-2 N· m-2= Pa Density ρ = m / V kilogram per cubic meter kg· m-3

Since antiquity there is the idea that things in nature are composed of very small particles, indivisible, called atoms (atom, gr. = indivisible). Atomic conception was resumed by J. Dalton in 1805, who showed on the basis of experimental observations accumulated until then that matter is made ​​of atoms. Avogadro's research related to gaseous state, led to the discovery of molecules, where the atoms are bound together in different ratios.

Atom = basic unit of matter, indestructible by ordinary chemical methods. Molecule = the smallest particle of a substance that can exist in the free state and manifests the properties of the substance.

At the end of the 19th century, the English physicist J. J. Thomson showed that the atoms of any element emit tiny negative particles. He called the particles "corpuscles", but later scientists preferred the name electron. Although atoms contain negative particles, the whole atom is not charged. So, Thomson concluded that atoms must also contain positive particles to balance the negative charge of the electrons. He imagined the “plum pudding model” of the atom, in which electrons are scattered like plums into the uniform “soup” or “cloud” of positive charges.

Plum pudding model electrons cloud of positive charge In this model, the atom is composed of electrons surrounded by a “soup” of positive charge to balance the electrons' negative charges.

19 The “plum pudding” model was disapproved by RUTHERFORD’s experience.
I. STRUCTURE OF SUBSTANCES The “plum pudding” model was disapproved by RUTHERFORD’s experience. Rutherford measured the deviation of alpha particles (helium ions with a positive charge) directed normally onto a sheet of very thin gold foil. Since the mass of the electrons is very small, and the α particles are heavy, considering the plum pudding model, the alpha particles should pass through the foil without significant deviation. The results of the experiment were very different from what Rutherford anticipated. Most of the α particles passed straight through the foil, but some of them were deflected at large angles and some were even reflected backward.

20 Rutherford’s experiment
Observed results Expected results for the “plum pudding model” “It was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you. … I realized that this scattering backward must be the result of a single collision, and when I made calculations I saw that it was impossible to get anything of that order of magnitude unless you took a system in which the greater part of the mass of the atom was concentrated in a minute nucleus. It was then that I had the idea of an atom with a minute massive center, carrying a charge” — Ernest Rutherford

Conclusions: the plum pudding model for the atom could not be correct; the atom must contain a very small (compared with the size of the atom) positive charge which causes the large deflections of the α particles; the atom is mostly empty space because most of the α particles passed directly through the foil.

Based on these results, Rutherford proposed the planetary model of the atom. According to this model the atom is composed of 2 parts: a small, compact center of positive charge called nucleus surrounded by a cloud of tiny negative particles called electrons, moving around the nucleus.

Planetary model of the atom

The planetary model of the atom is not perfect, because it can not explain why electrons are not colliding with the nucleus. It is known that an accelerating electric charge emits electromagnetic waves, so electrons will continuously lose energy, they will spiral towards the nucleus and finally they will collide with the nucleus. To explain why this did not occur, Niels Bohr (1913) postulated:

Electrons can move around the nucleus only on certain orbits, called allowed orbits. Electrons can gain or lose energy only by jumping from one allowed orbit to another. When an electron moves towards the nucleus energy is radiated and if it moves away from the nucleus energy is absorbed. For an electron to remain in its orbit the electrostatic attraction force between the electron and the nucleus must equal to the centrifugal force which tends to throw the electron out of its orbit.

Niels Bohr admitted that the orbits of the electrons are circular. Bohr’s model of the atom

A. Sommerfeld (1916), based on the atomic spectra of hydrogen, suggested that the permissive orbits of the electrons may be elliptic. Bohr – Sommerfeld model of the atom

In 1919 Rutherford discovered that the nucleus contains positive particles named protons. In 1932 James Chadwick discovered that nucleus contains also neutral particles – neutrons. Protons and neutrons are known as nucleons. + Nucleus Proton Neutron Nucleons

29 Properties of the main subatomic particles
I. STRUCTURE OF SUBSTANCES Properties of the main subatomic particles Electric charge Mass Symbol [C] Relative charge kg amu Proton +1.602·10-19 +1 ·10-27 1.0073 p; p+ Neutron ·10-27 1.0087 n; no Electron -1.602·10-19 -1 ·10-31 e; e- Proton = a nucleon with positive charge (+1) Neutron = a nucleon with zero charge (0) Protons and neutrons have close weights, and electrons are much lighter.

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