The Bear Bones PDF - Vanguard Science Department

THE BEAR BONES By The Vanguard Science Department (VSD) Introduction Within the confines of these pages lies a formidable guide to the SE course. The following is a warning to all that read the material within this document. Some chapters possess the power to lull you into a slumber enviab...

THE BEAR BONES By The Vanguard Science Department (VSD) Introduction Within the confines of these pages lies a formidable guide to the SE course. The following is a warning to all that read the material within this document. Some chapters possess the power to lull you into a slumber enviable even to Sleeping Beauty. The author can affirm that drowsiness does not only affect the reader, but at times, the writer. If you doubt the author, try writing your own set of notes on the relevant material. Readers will wade through explanations of matter where grams, without a preceding insta, affect the outcome of triangular equations. The ever-present demon of abject boredom will rear out of these pages. The mind-altering master Toc-of-tic will whisper sweet meaningless content. While the lord of streaks, Chats-of-snap, will remind you of a way to contact friends who care for the sake of their streak game. Beware of the real Bee that will briefly distract you but can lead to the lord of procrastination, Tubes of You. As for the older reader, the book of faces will be calling to you, but the author is sure you can control the desire to comment and inform others about how great your life really is. Throughout this seemingly endless tome, readers are likely to experience confusion, frustration, and at some point, a vague sense of recollection. Beware! Some experiences can summon the darkest creatures, known as “Why”, “What's the point,” and “Who cares”. "Why" will force you to reread the material or seek other sources of information. "What’s the point," the author will remind you of the dual objectives: to triumph in your exams and to cultivate knowledge, empowering you to explore the world. "Who cares," the author will assert, that the author indeed cares, despite their denial (but really, the author does care). It must also be pointed out that from chapter 1 onward there will be no attempts at humor. Mistakes are simply mistakes (please inform the author if any are found), and laughter is strongly discouraged, as it may bring an unwanted sense of enjoyment. Here begins your adventure. The author (who really does care) Course outline This course is split into 5 chapters. The first 3 chapters look into matter in some detail (Elements, Compounds, Mixtures). The fourth explores the effect of force and energy (basically really world math). The final chapter will look into some aspects of environmental science which are not covered in the AST course. Matter Everything in the visible universe that has a mass and volume is considered matter. Below is a diagram showing a way to separate matter into its parts. Content 1. The Building Blocks 1.1. Elementary particles 1.2. History of the atomic model 1.3. The sub atomic particles 1.4. The Simplified atomic model 1.5. The periodic table 1.6. Ions 1.7. Isotopes 1.8. Molecules and compounds 1.9. Balancing chemical equations 1.10. Ionic bond 1.11. Covalent bonds 1.12. Atomic model (Ball and stick - Lewis Notation)  The Building Blocks 1.1 Elementary particles In previous years we have looked at Atoms and Elements If you look inside an atom you find 3 subatomic particles which make up the atom. Just How Small is an Atom? The sub atomic particle gives an atom its mass and its charge. The mass of an atom can be found in the nucleus. The charge comes from a balance between the positively charged protons and the negatively charged electrons. The table below provide information about the subatomic particles Particle Location Relative Mass Charge Proton Nucleus 1 +1 Neutron Nucleus 1 Neutral Electron Electron Shell 0 -1 1.2 The History of the atomic model The history of the atomic model started over 2400 years ago in ancient Greece. The 2,400-year search for the atom - Theresa Doud 400 BC - Democritus Ancient Greek philosophers Democritus and Leucippus, suggested that all matter is composed of indivisible and indestructible particles called atoms. However the rest of the Greeks decided (they wanted to watch Avatar - Not true) that the world was made of four elements: earth, air, fire and water. 1803 - John Dalton In the early 19th century, John Dalton formulated the first modern atomic theory, proposing that elements are made up of indivisible atoms, and compounds are formed by combining atoms in simple, whole-number ratios. John Dalton came up with the idea that atoms were balls of different masses 1. All matter is made up of atoms 2. All the atoms of a particular element are the same (same mass, same size, same chemical properties) 3. One elements atoms are different from other elements atoms 1897 - J. J. Thomson J.J. Thomson discovered the electron and proposed the "plum pudding" model, suggesting that atoms were a positively charged mass with negatively charged electrons embedded within it like plums in a pudding. 1909 - Rutherford Ernest Rutherford conducted the famous gold foil experiment, leading to the discovery of the atomic nucleus and proposing a nuclear model of the atom. In this model, electrons orbit a small, dense, positively charged nucleus at the center of the atom. 1913 - Rutherford-Bohr Niels Bohr enhanced the atomic model by incorporating principles of quantum mechanics. He suggested that electrons orbit the nucleus in quantized energy levels, and electrons can absorb or emit energy in discrete packets or quanta. 1932 - Chadwick Chadwick's groundbreaking experiments, conducted in 1932, involved bombarding a thin sheet of beryllium with alpha particles (helium nuclei). The beryllium emitted a previously unidentified radiation that was not affected by electric or magnetic fields, indicating it consisted of electrically neutral particles. Chadwick concluded that these neutral particles had a mass similar to that of a proton but without an electric charge. He named these particles "neutrons." 1.3 The Simplified atomic model The simplified atomic model shows: - The number of protons, neutron - The number of electrons - The electron energy level (Orbit, Electron shell) The examples below show the two different methods of showing the atomic model. 1.4 Subatomic particles The atom below shows an atom's structure; the table explains the subatomic particle's location. Particle Symbol Location Mass (g) Relative Mass Charge (u) Proton 𝑃 + Nucleus 1. 673 × 10 −24 1. 007 ∼ 1 +1 Neutron 𝑛 Nucleus 1. 675 × 10 −24 1. 008 ∼ 1 Neutral Electron 𝑒 − Electron 9. 109 × 10 −28 0. 00055 ∼ 0 -1 Shell Charge The electronic charge of an atom is determined by the number of electrons (negative charge) and protons (positive charge). A neutral atom such as the one above has the same number of electrons as protons. If there is an unbalance the atoms is referred to as an Ion (see the section on Ions) Mass The mass of an atom comes for the number of protons and neutrons (See Isotopes) they are located in the nucleus at the centre of the atom The Organisation of the Electron Shells The table below shows the distribution of electrons through the electron’s energy levels. The first level (shell) must be filled before the second, the second before the third and the third before the fourth. Energy level Maximum number of electrons First 2 Second 8 Third 8 Fourth 2 (the table above is not strictly correct but it’s what we use in highschool) In reality the third shell can hold up to 18 electrons. However, once the third shell has 8 electrons the fourth shell begins to fill. In this way, the number of valence electrons always adds up the element group on the periodic table (also explaining why we only look at the first 20 elements this year) The diagram below shows the electron rings for an atom with 20 electrons 1.5 The periodic table The periodic table is a collection of all the known elements. 1. The periodic table is ordered by the Atomic number (Protons) 2. Each element is shown in a box (see below) The guide to reading the periodic table is shown below. The periodic - table a closer look To get a closer understanding of the periodic table, we are going to look at the diagram below. Pink (the size of the atom) Group 1 atoms have one valence electron and Group 8 atoms have a full outer shell of electrons. Each row on the periodic table is another ring therefore the atoms get bigger as you move down the periodic table. Group 1 one pull on there valence electron is less than those in group 8 This means that group 1 atom’s valence electrons is pulled away from the nucleus by other nucleus making it bigger. Orange (ionization energy) Ionization energy is the energy needed to rip electrons away from the nucleus and it works oppositely from the size of the atom. The explanation is that the smaller the atom, the more energy is needed to pull electrons away from the nucleus and it becomes easier for atoms to gain electrons instead. The ionization energy increases as you go down a row (left to right), but decreases as you go down the column because there are more electron shells and the valence electrons are further from the protons. Blue (Mass) More protons means more weight Green (Number of oxidation states) More oxidation states because there are more valence electrons that can be taken away. Iron can be oxidized into Iron (III) or Iron (II) because there are different numbers of electrons. Red (How metal is the element) That’s where the metals are on the periodic table. 1.6 Relative atomic mass and Isotopes An isotope is an atom with the same number of protons and a different number of neutrons. An atom of Carbon - 12 has 6 protons and 6 neutrons which means it will have an atomic mass of 12. If we find Carbon (6 Protons) but with a different number of neutrons as shown in the table below the same element will have different masses. Carbon 12 6 protons 6 Neutrons Mass of 12 Carbon 13 6 protons 7 Neutrons Mass of 13 Carbon 14 6 protons 8 Neutrons Mass of 14 Isotopes and the atomic mass of elements You will notice when looking at the periodic table that the atomic mass is not a whole number. This is due to two reasons, the first reason is that electrons have a tiny amount of mass approximately 0.0005u and the proton and neutron weigh slightly more than 1u. The second and main reason is that isotopes happen naturally and the atomic mass is an average of the element isotopes. Example 1 The table below shows the percentages of carbon isotopes found in nature. Isotope Mass Percentage Carbon 12 12 98% Carbon 13 13 1.5% Carbon 14 14 0.5% To find the atomic mass, we find the average of the masses based on the prevalence of each isotope Atomic mass = (% I1 x Mass I1) + (% I2 x Mass I2) + (% I3 x Mass I3) + etc….. Atomic mass = (0.98*12) + 0.015*13) + (0.005*14) Atomic mass = 12.025 u Example 2 A sample of carbon has been found. It is a mixture of Carbon 12 and 13. The percentage of these atoms is shown in the table below identifying the atomic mass of the sample. Isotope Mass Percentage Carbon 12 12g 70% Carbon 13 13g 30% To find the atomic mass, we find the average of the masses based on the prevalence of each isotope Atomic mass = (% I1 x Mass I1) + (% I2 x Mass I2) + (% I3 x Mass I3) + etc….. Atomic mass = (0.7*12g) + (0.3*13g) Atomic mass = 12.3 g 1.7 Ions Atoms become ions by losing or gaining valence electrons. Atoms lose and gain electrons because they want their outer ring to be full. An atom with a positive charge (more proton than electrons) is called a cation. An atom with a negative charge (less proton than electrons) is called an anion. Example A neutral atom of carbon has 6 protons and 6 electrons Gain 2 electrons 6 protons 8 electrons Carbon 2- ion Lose 1 electrons 6 protons 5 electron Cardon 1+ ion Cations and anions are fundamental concepts in chemistry related to the charge of ions, which are atoms or molecules that have gained or lost electrons. Cation: A cation is an ion with a positive charge. ​ ​ This positive charge is the result of the loss of one or more electrons from the atom's outer shell. When an atom loses electrons, it becomes positively charged because there are more protons (positively charged particles) than electrons (negatively charged particles). Cations are typically formed by metals and are attracted to negatively charged electrodes during electrolysis. Example: Sodium (Na) can form a cation by losing one electron: Na⁺ Anion: An anion is an ion with a negative charge. ​ This negative charge is the result of gaining one or more electrons in the atom's outer shell. When an atom gains electrons, it becomes negatively charged because there are more electrons than protons. Anions are typically formed by non-metals and are attracted to positively charged electrodes during electrolysis. Example: Chlorine (Cl) can form an anion by gaining one electron: Cl⁻ In summary, cations have a positive charge because they have lost electrons, and anions have a negative charge because they have gained electrons. These charges are crucial in understanding chemical reactions, ionic compounds, and the behavior of elements in various chemical processes. The diagram below shows how cations and anions can be formed. In the diagram an Sodium (Na) atom with one valence electron loses the electron to a Fluorine (F) atom with 7 valence electrons. The movement of electrons will make a Sodium +1 cation and a Fluorine -1 Anion. The most common ions Groups 1, 2 and 3 (13) atoms commonly lose electrons becoming cations. Groups 5 (15), 6 (16) and 7 (17) atoms more commonly gain electrons becoming Anions Group 4 (14) atoms both lose and gain electrons 1.8 Molecules and compounds A molecule is two or more atoms bound together. A compound is two or more different molecules bound together. How and bonds formed The formation of these bonds is driven by the desire of atoms to achieve a more stable and lower energy state. In most cases, this involves achieving a full outer electron shell or a stable electron configuration similar to that of noble gases. The specific type of bond formed depends on the properties of the atoms involved, including their electronegativity and the number of electrons in their outermost shell. Example of these are shown in the table below Molecule Atoms Type of bonding Diagram Table Salt Sodium and Ionic chlorine Water Oxygen and 2 Covalent hydrogen atoms Oxygen Two Oxygen Covalent (double atoms bond) 1.9 Ionic bond Ionic bonding is between a metal and a nonmetal. Ionic bonding involves the transfer of electrons from a metal to a non-metal, leading to the formation of ions that are attracted to each other, resulting in the creation of an ionic compound. Ionic compounds often have high melting and boiling points and are good conductors of electricity when dissolved in water or melted. Why does this happen? 1. An electron is most stable when the outer orbit is complete, this is done by transferring (giving or taking) electrons. 2. Metals normally want to give away electrons as it requires less energy to give electrons away when there is a low number of valence electrons. 3. Nonmetals want to gain electrons as it requires less energy for them to gain electrons to fill their outer orbit. 4. The movement of electrons then creates cations and anions. The electrostatic attraction between oppositely charged ions results in the formation of an ionic bond Example of Ionic Bond 1. The electron from the sodium (Na) feels the pull from Chlorine (Cl). 2. The pull from chlorine is stronger than the pull from sodium. 3. The electron leaves sodium and joins the chlorine atoms. 4. The movement of the electron creates two ions. One positive and one negative. 5. The positive and negative ions attract each other and form the ionic compound called sodium chloride (table salt). 1.10 Covalent bonds A Covalent bonds is between two nonmetals Covalent bonds involve the sharing of electrons between atoms, and they are common in molecular compounds composed of non-metal elements. The sharing of electrons helps both atoms achieve a more stable electron configuration, leading to the formation of molecules with distinct structures and properties. Why does this happen? In a covalent bond the atoms share electrons to fill the electron shells. Example: Methane 1. Methane is a combination of 4 Hydrogen atoms and one Carbon atom. 2. Once again, the atoms are trying to get a full outer shell. 3. By sharing electrons, both the hydrogen atoms and the carbon atom have a full outer shell. 4. As the electrons are being shared, it creates a bond between the atoms. 1.11 Atomic model Ball and stick A ball and stick diagram is a method used to show the number of bonds and the atoms in a molecule Group in the periodic table 1 2 3 4 5 6 7 8 Most likely charged state +1 +2 +3 +/- -3 -2 -1 Non 4 number of ionic bonds 1 2 3 4 3 2 1 0 number of covalent bonds 4 3 2 1 0 You will notice in the table that the number of bonds is linked with the amount of space in the outer shell and the number of valence electrons. Ionic bond 1. The example shows a molecule composed of Fluorine and Boron (𝐵𝐹3) 2. Boron ‘B’ (which is in group 3) has a charge of -3 and can form 3 ionic bonds. 3. Florine ‘F’ (which is in group 7) has a charge of +1 and can form 1 ionic bond. 4. To make a neutral atom, the ions must add up to 0. 5. So 1 Boron atom needs 3 Fluorine atoms, as shown in the diagram below. (-3 + 3*(+1) = 0) Covalent bond In a covalent bond the bonds are formed by sharing the electron 1. The example show a molecule of Carbon dioxide 𝐶𝑂2 2. Carbon ‘C’ (which is in group 4) has 4 valence electrons 3. Oxygen ‘O’ (which is in group 6) has 6 valence electrons. 4. The number of covalent bonds an atom forms is the number of additional electrons it needs to form a full outer ring. Lewis Notation The Lewis notation method of showing atoms is shown in the diagram below. - The Element symbol is in the center and the valence electrons are shown on the outside. The Lewis notation method Group in the 1 2 3 4 5 6 7 8 periodic table Element Lithi Magne Bor Carbo Nitrog Oxyg Florine Argo um sium on n en en n Valence 1 2 3 4 5 6 7 8 electron Lewis Notation Ionic bond The lewis dot method to show an ionic bond between Boron and Fluorine: Covalent bond The lewis dot method for showing a covalent bond between carbon and oxygen. Notice how the electrons are still present on their respective element, but they are now being shared.