Grade 10 Science notes lecture PDF

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Grade 10 Science notes lecture/class 1 Unit 1: Matter and Energy in Chemical Change Lesson 1: Alchemy and Atomic Models Key Points: 1. Understand how theories of the structure of the atom were developed. 2. Learn the parts of the atom and their characteristics. Discovery and Development of the Atomic Theory Part I: Alchemy and the birth of Chemistry Long before the science of chemistry existed, people made use of chemical reactions to dye cloth, tan leather, and prepare foods. Eventually people began to search for explanations for the structure and behavior of matter. One of the earliest “Atomists” was the Greek philosopher Democritus, he proposed that matter was made of indivisible particles called atoms. His ideas were overshadowed however by a more famous philosopher, Aristotle, who proposed that all matter was made of four elements, earth, air, water and fire. As a result the idea of atoms was not discussed seriously for 2000 years. During the middle ages chemistry reappeared because of the rise of a group called the alchemists, whose goal was to change common substances into gold. Although they were never successful in this quest they did make many important discoveries that helped to develop science. They produced many experimental procedures and laboratory apparatus. Their experiments yielded a wealth of knowledge about the characteristics of substances. Later the contributions of Roger Bacon and Antione Lavoisier changed chemistry from a trial and error process of observation to a science of measurement, and paved the way for modern chemistry and the atomic theory. Suppose we take a small cube of the element lead and cut it into smaller and smaller pieces. As the soft grey metallic pieces get smaller and smaller they still retain the properties of lead. Eventually we would reach a point when the particle of lead could no longer be divided and still retain its properties. This particle is an atom the smallest particle of an element that retains the properties of that element. • It was not until the late 1700s that chemists were able to relate chemical changes to events at the level of individual atoms. At that time the English chemist John Dalton first stated his atomic theory. Dalton’s Atomic Theory included the following ideas: 1. All elements are composed of tiny indivisible particles called atoms. 2. Atoms of the same element are identical. The atoms of any one element

are different from those of any other element. 3. Atoms of different elements can combine with one another in simple whole number ratios to form compounds. 4. Chemical reactions occur when atoms are separated, joined, or rearranged. However atoms of one element are not changed into atoms of another by a chemical reaction.

This theory has been generally accepted over the years although we now know that an atom can be broken down into smaller particles. Dalton believed in the “solid sphere model” which said that an atom is nothing more than an infinitely small solid sphere. Later in 1897 J.J. Thomson discovered that atoms contained small negatively charged particles called electrons. So he modified Dalton’s model into what he called the “Plum Pudding model” which said that an atom was a solid sphere with small negatively charged particles embedded in it like plum pudding had pieces of plum in it. Shortly thereafter, Protons were discovered, this was due to the fact that all chemists agreed that atoms had to be electrically neutral, so if there was a small negatively charged particle there had to be a positively charged particle as well, this positively charged particle is many times larger than the electron and is called the Proton. In 1932 James Chadwick discovered that Protons were not the only large particle in the atom. He found that the mass of many elements was too large to be accounted for by only protons and electrons. His experiments proved the existence of the third and final subatomic particle the Neutron which carries no charge or is neutral. Further to this discovery a scientist named Ernest Rutherford felt that atoms could not have all these particles floating around randomly, and that the plum pudding and solid sphere models were not accurately explaining the findings of later experiments. So he conducted an experiment where he bombarded a thin sheet of gold foil with alpha particles. If the protons, neutrons and electrons were simply floating about (which was the current theory) then the alpha particles should pass through the foil unhindered. To the shock of Rutherford and his associates many of the alpha particles bounced back. Rutherford was so impressed that he was quoted as saying “it was like firing a 15inch shell at a piece of tissue paper and having it bounce back at you!”. This led to the development of the Rutherford model of the atom which said the atom was made of a dense central nucleus where the heavy particles (protons and neutrons) were located and a the electrons surrounded this nucleus. Rutherford then developed this model of the atom (at left).

Later experiments by Neils Bohr showed this to be mostly correct except that the electrons are not positioned randomly they can be found orbiting the nucleus at fixed distances or energy levels. Further to this Louis de Broglie surmised that electrons could not be found in fixed positions, but did show evidence of distinct energy levels, so he combined the Rutherford nuclear model and the Bohr model into what is known as the “electron cloud model”. In de Broglie’s model the electrons can jump or fall from one level to another depending on the energy of the atom. This is what allows for “glow in the dark” fluorescent type materials, electrons in the atom gain energy from light and are excited to a new level. When the light is turned off they slowly return to their previous state, releasing their energy in the form of light. The de Broglie model is currently the accepted model. The

structure of the atom, as we currently understand it is that there is a dense positively charged nucleus that contains the neutrons and protons, and there is a large volume of space surrounding the nucleus in which the electrons can be found. So in fact atoms are made up of a lot of empty space! So if all atoms are made of protons neutrons and electrons, how can the atoms of one element differ from another? The answer is that all elements have different numbers of protons in their nucleus, Carbon atoms, for example, have 6 protons while lead atoms have 82 protons, and fluorine atoms have 9 protons. These differing numbers of subatomic particles influence the properties of the different elements. Lesson 1.1: Isotopes, Electron Configurations Atomic structure and the Periodic Table Key Points: 1. Learn what Isotopes are and their importance. 2. Learn and understand the ideas of Bohr on the structure of the atom. 3. Follow the progression from Bohr’s model to the Quantum Mechanical (Electron Cloud) model of the atom. 4. Understand how the periodic table reflects these ideas. Isotopes Dalton stated that all atoms of each element are identical. We know now that while this is essentially true (number of protons in each element’s atoms is the same), it is not entirely true (some atoms of the same element have different numbers of neutrons) and are called isotopes. While these different atoms are chemically alike (because the number of particles that determine chemical properties – protons and electrons – are the same), they do have different masses. To show the differences in the Isotopes we

write the chemical symbol for that element with two different numbers to its left. The mass number is the superscript (top number) and the atomic number is the subscript (lower number). Atoms are then named based on their mass number i.e. Helium 4 or Uranium 238. Atomic Mass The atomic mass of an element then reflects the weighted average of all the isotopes of each element. This means the mass of each isotope is averaged together based on its abundance in nature: As an example, Sulfur has 4 isotopes: Sulfur 32 has an atomic mass of 31.972 and comprises 95% of all natural samples, while Sulfur 33 has an atomic mass of 32.971 but makes up only 0.76% of all natural samples. Sulfur 34 has a mass of 33.967 making up 4.22% of natural samples, and finally Sulfur 36 has a mass of 35.967 but accounts for a miniscule 0.014% of natural samples. Thus the atomic mass of Sulfur is calculated as by multiplying each atomic mass by the percentage abundance and then adding them all together. (31.972x0.95) + (32.971x0.0076) + (33.967x0.0422) + (35.967x0.00014) = 32.06 The Bohring World of Niels Bohr In 1913 Bohr proposed that electrons are arranged in concentric circular paths or orbits around the nucleus. Bohr answered in a novel way why electrons which are attracted to protons, never crash into the nucleus. He proposed that electrons in a particular path have a fixed energy. Thus they do not lose energy and crash into the nucleus. The energy level of an electron is the region around the nucleus where it is likely to be. These energy levels are like rungs on a ladder, lower levels have less energy and electrons must gain energy to move to a higher level in the same way that to climb a ladder you do work. The opposite is also true if an electron loses energy it falls to a lower level. Also an electron can only be found on a path, not between them in the same way that you cannot stand in between the rungs of a ladder. The amount of energy gained or lost by every electron is not always the same. Unlike the rungs of a ladder, the energy levels are not evenly spaced. A quantum of energy is the amount of energy needed to move an electron from its present energy level to the next higher one or to make a quantum leap.

The Quantum Mechanical Model Like the Bohr model, the quantum mechanical model leads to quantized energy levels for an electron. However the Quantum Mechanical model does not define the exact path an electron takes around the nucleus. It is concerned with the likelihood of finding an electron in a certain position. This probability can be portrayed as a...