Atoms and Atomic Theory PDF

CHAPTER 2: Atoms and atomic theory 1 Index 2-1 First Discoveries in Chemistry and Atomic Theory 2-2 Electrons and Other Discoveries in Atomic Physics 2-3 Atomic Nuclei 2-4 Chemical Elements 2-5 Atomic Mass 2-6 Periodic Table 2-7 Mole Conce...

CHAPTER 2: Atoms and atomic theory 1 Index 2-1 First Discoveries in Chemistry and Atomic Theory 2-2 Electrons and Other Discoveries in Atomic Physics 2-3 Atomic Nuclei 2-4 Chemical Elements 2-5 Atomic Mass 2-6 Periodic Table 2-7 Mole Concept and Avogadro's Number 2-8 Using the Mole Concept in Calculations 2 1  2-1 First Discoveries in Chemistry and Atomic Theory Antoine Lavoisier (1774) Law of Conservation of Mass The sum of the masses of the products resulting from the reaction; is equal to the sum of the masses of the reacting substances. LAVOISIER Joseph Proust (1799) Law of Constant PROUST Proportions DALTON All samples of a compound have the same THOMPSON composition. The components combine MILLIKEN at a constant rate by mass. RUTHERFORD BOHR SCHRÖDİNGER 3 Dalton's Atomic Theory 1. Each element is composed indivisible particles (called atoms). Atoms are neither created nor destroyed in chemical reactions. 2. All atoms of a given element are alike in mass (weight) and other properties but differ from all other elements 3. A chemical compound is formed by the combination of two or more elements in a simple numerical ratio. 1766-1844 British scientist For example, one atom A and one atom B (AB) or one atom A John Dalton is the scientist who and two atoms B (AB2). put forward the first atomic Dalton's atomic theory also enabled us to understand «The Law theory. of Multiple Proportions». 4 2  Consequences of Dalton's Atomic Theory  In the formation of carbon monoxide (CO), 1.33 g of oxygen combines with 1.0 g of carbon.  In the formation of carbon dioxide (CO2), 2.66 g of oxygen combines with 1.0 g of carbon. The second oxide is richer in oxygen and contains twice as much oxygen as the first. 2,66 g / 1,33 g = 2,00 CO and CO2 1:1 and 1:2 5 2-2 Electrons and Other Discoveries in Atomic Physics Forces Between Electrically Charged Objects (b) Since objects do not carry any electrical charge, they do not exert force on each other. (c) Objects carry opposite charges (one positive and (a) Objects carrying similar the other negative) and charges repel each other. attract each other. 6 3  Coulomb's law (1785): The magnitude of the attractive electrostatic force between two point charges is directly proportional to the product of the magnitudes of their charges and inversely proportional to the squared distance between them. 7 Coulomb's law 1. The direction of the electric field vector points outward for positive charges and inward for negative charges. 2. It creates an electric field around the charge. 3. The magnitude of the electric field is proportional to the charge. 8 4  Effect of Magnetic Field on Charged Particles If charged particles are sent perpendicular to a magnetic field, they are deflected by the field. Negatively charged particles are deflected in one direction, positively charged particles are deflected in the opposite direction. 9 Discovery of Electrons Cathode Ray Tube (CRT): The first cathode ray tube was made by Michael Faraday (1791- 1867) approximately 150 years ago. Faraday discovered cathode rays emitted by the negative end (cathode) when he passed an electric current through a glass tube that was largely evacuated of air. 10 5  Cathode Ray Tube The high voltage source creates a negative charge on the left electrode (cathode). It creates a positive charge on the right electrode (anode). In reality, cathode rays are not visible to the naked eye, but they can be seen by the light they emit from a surface they hit. 11 Thomson proposed that cathode rays were negatively charged elementary particles found in all atoms. From now on, cathode rays were called electrons. The first person to use this term was George Stoney (1874). 12 6  Charge of the Electron  Robert Millikan (1868-1940) determined the electron charge (e) by performing a series of 'oil drop' experiments between 1906 and 1914.  Millikan determined the electron charge. The charge of the electron is 1.6022x10-19 C.  If this value is substituted in the mass/charge relation of the electron, the value for mass (m) is 9.1094 x 10-28 g. 13 According to the J.J Thomson atomic model, a neutral atom should have positive charges that balance the negative charge, and these positive charges should be in the form of a cloud. Electrons must float in this cloud of positive charge. Pozitive charged nuclei Negative charged electrons Thomson Atom Modeli 14 7  RADIOACTIVITY 15 X-rays and Radioactivity Radioactivity: It is the phenomenon in which an unstable atomic nucleus gives off small matter particles and electromagnetic radiation. In 1895, while Wilhelm Roentgen (1845-1923) was working with cathode rays, he noticed that some substances outside the cathode ray tube were glowing or emitting fluorescent light. Since the nature of this light was unknown, Roentgen called it X-ray. Ernest Rutherford discovered two of the rays emitted by radioactive substances, alpha (α) ray and beta (β) ray. 16 8  Rutherford Atomic Model Rutherford explained the results of this experiment in 1911. Rutherford atomic model has the following features: 1. Most of the mass of an atom and all of its positive charge is concentrated in a very small region called the nucleus. Most of the atom is an empty piece of space. 2. The size of the positive charge varies from atom to atom and is approximately half the atomic weight of the element. 3. Outside the nucleus there are a number of electrons equal to the nuclear charge. The atom itself is neutral in electrical charge. 18 Atomic Nucleus - Discovery of Protons and Neutrons Rutherford discovered protons (1919) Here the electrons are drawn much closer to the nucleus than they actually are. In reality; If the atom is a room with dimensions of 5m x 5m x 5m, the nucleus is only as much as the space occupied by this written sentence. James Chadwick discovered neutrons (1932) 19 9  2.3-STRUCTURE OF THE ATOM Matter is made of atoms. The smallest part of a substance that cannot be broken down by chemical means is called an atom. The atom consists of 3 fundamental particles: Proton: It is located in the nucleus of An atom whose number of electrons is the atom and has a positive (+) charge. equal to the number of protons is called a neutral atom, and an object whose Neutron: It is found in the nucleus of atoms are neutral is called an the atom and has no charge. uncharged object. However, being called uncharged Electron: They orbit around the means that the (+) and (-) charges are nucleus and have a negative (-) charge. equal. This means there is no excess load on it. 20 STRUCTURE OF THE ATOM Positively charged object: When a neutral object loses electrons, an excess of (+) charge occurs. Objects with such (+) excess charge are called positively (+) charged objects. Negatively charged object: When a neutral object gains electrons, an excess of (-) charge occurs. Objects with such (-) excess charge are called negatively (-) charged objects. 21 10  The properties of 3 Basic particles Particle Mass Electricle Charge SI(kg) (amu) SI(Coulomb) Atom Electron 9,109 x 10-31 0,00054858 –1,602 x 10-19 –1 Proton 1,673 x 10-27 1,0073 +1,602 x 10-19 +1 Neutron 1,675 x 10-27 1,0087 0 0 The atomic mass unit (amu) is exactly 1\12 of the mass of carbon- 12. 1 amu = 1.66054 x 10-24 kg 22 2.4- Elements Each element has a name and symbol; A= Mass number Z = Atomic number neutron number = A-Z Charge of ion= Number of proton- Number of electron Same atomic number (Z), different mass number (A) Atoms that have ions are called isotopes. 23 11  2-5 Atomic Mass The atomic mass (weight) of an element is the average of the weighted atomic masses of the isotopes according to their occurrence in nature. The weighted atomic mass is calculated according to the following general equation; Atomic mass of the element= 24 25 12 26 2-7 Mole Concept and Avogadro's Number The amount of an uncountable amount of a single type of substance is expressed in units of moles in the SI unit system. The number of particles contained in one mole of particles such as atoms and molecules is called Avogadro's number (NA). NA = 6,02214199 x 1023 mol-1 The mass of one mole of atoms is called molar mass, MA 27 13  2.8- Concept of Mole in Calculations Potassium-40 is one of the few isotopes of elements with small atomic numbers that are naturally radioactive, and its abundance among K isotopes in nature is 0.012%. How many 40K atoms will you ingest when you drink 225 mL of milk containing 1.65 mg/mL potassium? 28 14