Yearly Year 10 Science 2024 Assessment Notes PDF

Yearly Year 10 Science 2024 Assessment Notes Table of Contents Yearly Year 10 Science 2024 Assessment Notes 1 Student Research Project 3 Types of Chemical Reactions...

Yearly Year 10 Science 2024 Assessment Notes Table of Contents Yearly Year 10 Science 2024 Assessment Notes 1 Student Research Project 3 Types of Chemical Reactions 6 Rules For Naming Compounds 11 Endothermic and Exothermic Reactions 12 Endothermic Reaction Example 12 Exothermic Reaction Example 12 Law of conservation of mass 14 Radioactivity 15 Photosynthesis 16 Precipitation Reactions 16 Decomposition Reactions 16 The three types of decomposition reactions 17 Video Links 17 Part B 20 Types of Chemical Reactions 20 Atoms: 20 Elements: 21 Compounds: 21 Common Names vs. Chemical Formulae: 22 Range of Compounds: 22 Identifying Compounds: 22 Ionic Bonds: 22 Covalent Bonds: 23 Differentiating Ionic and Covalent Bonds: 23 Types of Chemical Reactions: 24 Acids and Bases: 25 Reactions Involving Acids and Bases: 25 Importance of Acid-Base Reactions: 26 Law of Conservation of Mass: 26 Number of Atoms in Chemical Reactions: 26 Importance of Conservation of Mass: 27 Part C 27 Basic Structure of Atoms: 27 Subatomic Particles: 27 Energy Levels or Shells: 28 Electron Configuration: 28 Isotopes: 28 Importance of Atom Structure: 28 1. Dalton's Model (Solid Sphere Model): 29 2. J.J. Thomson's Model (Plum Pudding Model): 29 3. Rutherford's Model (Planetary Model): 30 4. Bohr's Model (Bohr-Rutherford Model): 30 5. Quantum Mechanical Model (Modern Model): 31 Importance of Different Models: 31 Radioactivity and Atoms 31 Radioactivity: 31  Radioactive Elements and Isotopes: 31 Half-life of Radioactive Isotopes: 32 Isotopes in Industry 32 Advantages and Disadvantages of Using Isotopes in Industries: 32 Applications of Half-Life and Carbon - 14 33 Living with Radiation 34 Outcome 1: Students can identify the use of Radiation 34 Outcome 2: Students can research and analyse the effects of using radiation in everyday life 34 Coordination and Disease 41 Your Body's Battle: Innate Immune Defence 60 1. What is the main purpose of your innate immune system? 60 2. What is the role of phagocytes like macrophages in the innate immune system? 60 3. Why does the body increase temperature during an inflammatory response? 60 4. What is the purpose of the increased permeability of blood vessels during inflammation? 60 5. When the body's local innate defences become overwhelmed, what happens? 61 6. What is the role of natural killer cells in the innate immune system? 61 Motion 74 Physics Behind Car Crashes and Car Safety Features 79 1. Newton's First Law of Motion (Law of Inertia) 79 2. Newton's Second Law of Motion (F = ma) 79 3. Newton's Third Law of Motion (Action and Reaction) 80 4. Conservation of Momentum 80 5. Impulse and Force 80 6. Energy Absorption and Dissipation 81 7. Friction and Road Conditions 81 8. Rollover Prevention 81 Summary 82 Waves and Light 82 Lesson 8: Transmitters and Receivers 98 Additional Questions: 104 Global Systems 105 Cyclones 106 Earthquakes 106 Volcanic Eruptions 106 Greenhouse Gases (GHGs) 107 Ozone Layer 107 Long-Term Effects of Climate Change 108 Origins of the Universe 109 1. Origin and Age of the Universe: 109 2. Models of the Universe: 109 3. Force of Gravity: 109 4. Major Features Contained in the Universe: 109 5. Size of and Distances between Structures: 110 7. Timeline of Space Study: 110 1. Geocentric Model (Ancient - 16th Century) 111 2. Heliocentric Model (16th Century) 112 3. Kepler's Laws of Planetary Motion (17th Century) 113 4. Newton's Universal Gravitation (1687) 114 5. Static Universe (Early 20th Century) 115 6. Expanding Universe (1920s) 116 7. Big Bang Theory (Mid-20th Century) 117 8. Steady State Theory (Mid-20th Century, Discredited) 118 9. Modern Cosmology (Late 20th Century – Present) 119 Conclusion 120  Student Research Project Hypothesis: A hypothesis is a proposed explanation for a phenomenon. For a hypothesis to be a scientific hypothesis, the scientific method requires that one can test it. Scientists generally base scientific hypotheses on previous observations that cannot satisfactorily be explained with the available scientific theories. Aim: ‘To Investigate the Effect of….’ Equipment Hazard Safety Measure Method: Always start with a word of action such as ‘grab, lift, boil or stop’.   Types of Chemical Reactions Atom Smallest divisible part of matter , which consists of negatively charged electrons revolving around a positively charged nucleus, consists of protons and neutrons. Elements Purest form of a substance Compound Combination of two or more elements that are chemically bonded. Molecule Numbers of compound particles Chemical Reaction Process in which two or more chemical substances combine by rearranging their chemical bonds Reactants Substances that take part in chemical reactions Products Results that are produced from a chemical reaction.  Electronegativity Sodium Chloride is easier to take one electron which mean sodium get one negative charge Chlorine has 7 electrons in its outermost shell therefore the sodium(which has one in its outer shell) takes one electron from the chlorine to become an ionic compound. STP means standard temperature and pressure( Normal Temperature for it to exist. Also includes atmospheric pressure.) Covalent compounds can exist in all 3 states of matter. Metals are shiny, ductile and brittle Metal point of ionic compounds is very high, the electrons have a connection because of the positive and negative charge. To break the bond you need a lot of energy. Ionic bonds is the strongest because it has to opposite charged entities Covalent compounds do not conduct electricity Ionic bonds are formed when two atoms exchange electrons to create a positive and negative ion. Covalent bonds are formed when atoms share electrons to create a molecule. Metallic bonds are created when metal atoms lose their outermost electron to form positively charged ions. Covalent Compounds = A covalent bond is a chemical bond that involves the sharing of electrons to form electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs. The stable balance of attractive and repulsive forces between atoms, when they share electrons, is known as covalent bonding. Additionally a covalent compound is a molecule that is formed when two or more different atoms are connected with a covalent bond. Covalent compounds are basically the molecules that form when two different atoms form a covalent bond. They typically have similar electronegativity and are non-metals. Ionic Compounds = Ionic compound, any of a large group of chemical compounds consisting of oppositely charged ions, wherein electron transfer, or ionic bonding, holds the atoms together. Naming Compounds ➔ Sodium Chloride ➔ Sodium Hydroxide ➔ Potassium Nitrate ➔ Potassium Nitrite ➔ Carbon Monoxide ➔ Lead Sulphate  ➔ Sodium Dioxide Two ways Words Chemical formulas Compounds have constant composition with respect to mass because they are composed of atoms in fixed ratios. A chemical formula indicates the elements present in a compound and the number of atoms of each. Ionic Compounds Composed of positive and negative ions. Usually formed from metal and non-metal. 3 types: 1. Simple (Binary) Ionic compounds 2. Compounds with transition metals 3. PolyAtomic Ionic compounds Binary means 2 elements Ionic means a metal and a non-metal (or cation and anion) Writing Formulas from Names  1st word = CATION 2nd word = ANION name with ide ending A Compound is formed between potassium and iodine. Name the compound Potassium Iodide , KI The given elements are: One atom of barium Two atoms of chlorine Barium Chloride 𝐵𝑎𝐶𝑙2 Rules For Naming Compounds Type of reaction General equation Combination( synthesis) A+B→AB Decomposition AB→A+B Single Displacement A+BC→AC+B Double Displacement AB+CD→AD+BC Combustion 𝐶𝑥𝐻𝑦 + 𝑂2 −−> 𝐶𝑂2 + 𝐻2𝑂 Combination (Synthesis) 𝑍𝑛(𝑠) + 𝐶𝑙2(𝑔) −−> 𝑍𝑛𝐶𝑙2(𝑠) Decomposition 𝑁𝐻4𝐶𝑙(𝑠) −−> 𝑁𝐻3(𝑔) + 𝐻𝐶𝑙(𝑔) Single Displacement 𝑍𝑛(𝑠) + 𝐶𝑢𝐶𝑙2(𝑎𝑞) −−> 𝐶𝑢(𝑠) + 𝑍𝑛𝐶𝑙2(𝑎𝑞) Double Displacement 𝐴𝑔𝑁𝑂3(𝑎𝑞) + 𝐾𝐵𝑟(𝑎𝑞) −−> 𝐴𝑔𝐵𝑟(𝑠) + 𝐾𝑁𝑂3(𝑎𝑞) Combustion (Methane) 𝐶𝐻4(𝑔) + 2𝑂2(𝑔) −−> 𝐶𝑂2(𝑔) + 2𝐻2𝑂(𝑔) Endothermic and Exothermic Reactions Endothermic reactions are chemical reactions in which the reactants absorb heat energy from the surroundings to form products. An exothermic reaction is a reaction in which energy is released in the form of light or heat. The energy is absorbed from the surrounding into the reaction. Endothermic Reaction Example Melting ice, evaporation, cooking, gas molecules, and photosynthesis are a few examples. Exothermic Reaction Example Rusting iron, settling, chemical bonds, explosions, and nuclear fission are a few examples. Exothermic Reaction are more common Endothermic Reaction takes in heat Exothermic releases heat Problem: A sample of sodium hydroxide was dissolved in a test tube of water at 19 °C. The temperature increased to 28 °C. A sample of potassium nitrate was dissolved in a test tube of water at 20 °C. The final temperature was 11 °C. Determine the temperature change for each and state which reaction was endothermic and which was exothermic. Energy and Chemical Reactions Chemical Energy – Energy stored in the chemical bonds of a substance. Chemical reactions always involve energy changes. Making bonds and breaking bonds involve energy changes Activation Energy The energy required to break the bonds in the reactants for a chemical reaction to occur. Endothermic and Exothermic reactions Step 1: Energy must be SUPPLIED to break chemical bonds of reactants: Step 2: Energy is RELEASED when new chemical bonds are made in the products: Energy is SUPPLIED than is RELEASED then the reaction is ENDOTHERMIC Endothermic Reaction REACTANTS + ENERGY PRODUCTS OR REACTANTS + HEAT PRODUCTS Endothermic 6CO2 + 6H2O + Energy C6H12O6 + 6O2 Exothermic CH4 + 2O2 CO2 + 2H2O + Energy Exothermic Reactions REACTANTS PRODUCTS + ENERGY OR REACTANTS PRODUCTS + HEAT Law of conservation of mass Law of conservation of mass states that, during any chemical reaction, matter is neither created nor destroyed. Mass is conserved from reactants to products. Who Discovered this Law? 1789, France Antoine Lavoisier Scientist Used one of the first analytical mass balances to prove this law. Executed on the guillotine during the French Revolution. He is known as the “Father of Chemistry” because he made it a quantitative science. Therefore 𝑀𝐴𝑆𝑆𝑅𝐸𝐴𝐶𝑇𝐴𝑁𝑇𝑆 = 𝑀𝐴𝑆𝑆𝑃𝑅𝑂𝐷𝑈𝐶𝑇𝑆 “→” means “gives”. Small numbers (subscripts) – tell how many of a particular type of atom are inside of a molecule. Big numbers (coefficients) – tell how many of each particle is involved in the reaction.  Acid - a substance with particular chemical properties including turning litmus red, neutralising alkalis, and dissolving some metals; typically, a corrosive or sour-tasting liquid of this kind. + Acids have a high concentration of 𝐻 𝐼𝑜𝑛𝑠 Acids are below 7 on the PH scale Bases have a PH scale value greater than 7 The PH scale value of 7 is considered neutral Water has a PH scale value of 7 Base - A base is a substance that can neutralise the acid by reacting with hydrogen ions. Most bases are minerals that react with acids to form water and salts. The pH scale measures how acidic or alkaline a substance is. − Chemical substances having a pH greater than 7, have a high concentration of 𝑂𝐻 Radioactivity Radioactive decay is the process by which an unstable atomic nucleus loses energy by radiation. A material containing unstable nuclei is considered radioactive. When the atoms of an element have extra neutrons or protons it creates extra energy in the nucleus and causes the atom to become unbalanced or unstable. Three of the most common types of decay are alpha, beta, and gamma decay. Alpha decay is a nuclear decay process where an unstable nucleus changes to another element by shooting out a particle composed of two protons and two neutrons Beta Decay is a type of radioactive decay in which a proton is transformed into a neutron or vice versa inside the nucleus of the radioactive sample. Gamma decay, a type of radioactivity in which some unstable atomic nuclei dissipate excess energy by a spontaneous electromagnetic process. Photosynthesis Sunlight provides the energy needed for photosynthesis to take place. In this process carbon dioxide and water are converted into oxygen (a waste product that is released back into the air) and glucose (the source of energy for the plant). Plantsgo under photosynthesis to make their food, they use cellular respiration the energy from food to power the growth. Cellular respiration is the process in which cells break down molecules for energy in the form of ATP molecules. The cells use that oxygen that you breathe in and produce carbon dioxide that you breathe out. It is an exothermic reaction since it releases energy Respiration can be thought of as a combustion of glucose (without flames). It is very slow compared to other combustion reactions. Aerobic respiration is the process by which oxygen-breathing creatures turn fuel, such as fats and sugars, into energy. Anaerobic respiration transfers energy from glucose to cells. It occurs when oxygen is not present. It transfers large amounts of energy quickly. Precipitation Reactions A precipitation reaction refers to the formation of an insoluble salt when two solutions containing soluble salts are combined. The insoluble salt formed is known as the precipitate, hence the reaction’s name. Precipitation reactions occur when cations and anions in aqueous solution combine. When the two separate solutions of two soluble salts are mixed, ion pairings are swapped, and one or more insoluble salts are produced. The swapping of ions in the compounds is called displacement. The solids produced in precipitate reactions are crystalline solids, and can be suspended throughout the liquid or fall to the bottom of the solution. The remaining fluid is called supernatant liquid. The precipitate and supernatant can be separated by various methods, such as filtration, centrifuging, or decanting. A solubility table is used to determine if the products are soluble or insoluble. If they are insoluble they will be the precipitate. Decomposition Reactions Decomposition reactions are chemical reactions in which a more complex molecule breaks down to make simpler ones. The reaction can be generalised as: AB -> A + B  It is the opposite of synthesis reactions, where simpler reactants combine to form a more complex product. If a reaction has only one reactant, it is a decomposition reaction. A decomposition reaction is a type of chemical reaction in which one reactant yields two or more products. Most decomposition reactions are endothermic, since energy, either in the form of heat, electric current, or sunlight must be provided in order to break the bonds of the more complex molecule. The three types of decomposition reactions 1. Thermal decomposition reactions - energy in the form of heat is required to break the bonds of the more complex molecule. 2. Electrolytic decomposition reactions - such reactions occur when an electric current is passed through an aqueous solution of a compound. 3. Photo decomposition - these reactions occur in the presence of light (photons) Video Links (Some of this stuff covers Yr11 Chemistry so, just do what is in our Yr10 course) How To Name Ionic Compounds With Transition Metals - ✅ How To Name Ionic Compounds With Transition Metals Writing Ionic Formulas - Basic Introduction - ✅ Writing Ionic Formulas - Basic Introduction How To Name Acids - The Fast & Easy Way! - ✅ How To Name Acids - The Fast & Easy Way! Writing Chemical Formulas For Ionic Compounds - ✅ Writing Chemical Formulas For Ionic Compounds How To Name Ionic Compounds With Transition Metals - ✅ How To Name Ionic Compounds With Transition Metals How To Write Ionic Formulas With Polyatomic Ions - ✅ How To Write Ionic Formulas With Polyatomic Ions Types of Chemical Reaction (Decomposition) - ✅ Chemical Reactions - Combination, Decomposition, Combustion, Single & Double Displacement Chemistry Precipitation Reactions - ✅ Precipitation Reactions and Net Ionic Equations - Chemistry Endothermic and Exothermic Reactions - ✅ Endothermic and Exothermic Reactions Alpha Particles, Beta Particles, Gamma Rays, Positrons, Electrons, Protons, and Neutrons - ✅ Alpha Particles, Beta Particles, Gamma Rays, Positrons, Electrons, Protons, and Neutrons Acids and Bases - Basic Introduction - Chemistry - ✅ Acids and Bases - Basic Introduction - Chemistry Balancing Chemical Equations - ✅ Introduction to Balancing Chemical Equations Balancing Chemical Equations: UPDATED - Chemistry Tutorial Charges within the Periodic Table - Column 1 = +1 Charge. - Column 2= +2 Charge. - Column 3 = +3 Charge. - Column 18 = 0 Charge.  - The rest you will either have to remember or it will specify in certain questions in the test with roman numerals in brackets for example like = Iron (II) having at +2 charge in this case. - Protons - Neutrons = Valency of Element. Extra Test Revision Have an example of each of the following and definition = 1) Decomposition Reaction: A decomposition reaction is a reaction in which a compound breaks down into two or more simpler substances. The general form of a decomposition reaction is: AB→A+B. Most decomposition reactions require an input of energy in the form of heat, light, or electricity. Examples = Examples of decomposition reactions include the breakdown of hydrogen peroxide to water and oxygen, and the breakdown of water to hydrogen and oxygen. 2) Precipitation Reaction: Precipitation reactions occur when cations and anions in aqueous solution combine to form an insoluble ionic solid called a precipitate. Whether or not such a reaction occurs can be determined by using the solubility rules for common ionic solids. Examples = 1) 2KI(aq)+Pb(NO3)2(aq)→PbI2(s)+2KNO3(aq) 2) Pb2+(aq)+2I−(aq)→PbI2(s) 3) NaF(aq)+AgNO3(aq)→AgF(s)+NaNO3(aq)(molecular) 3) Acid Reaction: Examples = acid + metal → salt + hydrogen. acid + base → salt + water. acid + carbonate → salt + water + carbon dioxide. acid + hydrogen carbonate → salt + water + carbon dioxide. acid + ammonia → ammonium salt. 3 Main Acids in Tests = Hydrochloric Acid, Sulphuric Acid and Nitric Acid. 4) Single Displacement Reaction: A single replacement reaction, sometimes called a single displacement reaction, is a reaction in which one element is substituted for another element in a compound usually in the form A + BC = B + AC. Examples = 1. Mg ( s ) + Cu(NO 3 ) 2 ( a q ) → Mg(NO 3 ) 2 ( a q ) + Cu ( s ) 2. Al ( s ) + Fe 2 O 3 ( s ) ⟶ Δ Al 2 O 3 ( s ) + Fe ( s ) 3. Zn ( s ) + 2 HCl ( a q ) → ZnCl 2 ( a q ) + H 2 ( g ) 4. 2 Na ( s ) + 2 H 2 O ( l ) → 2 NaOH ( a q ) + H 2 ( g ) 5) Double Displacement Reaction: A double displacement is a reaction in which the positive and negative ions of two ionic compounds exchange places to form two new compounds. The double-displacement reaction generally takes the form of AB + CD → AD + CB where A and C are positively-charged cations, while B and D are negatively-charged anions. Examples = 1) Sodium oxide reacts with silver acetate. 2) Silver nitrate reacts with magnesium chloride.  3) Potassium carbonate reacts with ammonium iodide. 4) Cesium sulphide reacts with sodium hydroxide. 5) Sodium hydroxide reacts with hydrochloric acid. 1) Na2 O + 2AgC2 H3 O2 ⟶ 2NaC2 H3 O2 + Ag2O. 6) Combustion Reaction: A combustion reaction is a kind of chemical reaction in which a reaction between any combustible substance and an oxidiser takes place in order to form an oxidised product. Combustion reactions are often accompanied by fires and the release of energy in the form of heat. Examples: 1. Burning any kind of Wood or Coal to heat your home. 2. Cars and buses burn petrol or diesel to run. 3. Natural Gas or LPG is in use on your stovetop. 4. For the production of energy in thermal power plants. 5. Fireworks. 6. Combustion of butane (commonly found in lighters). Part B Types of Chemical Reactions Atoms: 1. Definition: Atoms are the basic units of matter. They are the smallest particles of an element that retain the chemical properties of that element. 2. Structure: Atoms consist of a nucleus containing protons (positively charged) and neutrons (neutral), surrounded by electrons (negatively charged) orbiting the nucleus in energy levels or shells. 3. Atomic Number: This is the number of protons in an atom's nucleus. It determines the element's identity. 4. Isotopes: Atoms of the same element with different numbers of neutrons are called isotopes. They have the same atomic number but different mass numbers. 5. Atomic Mass: The total mass of an atom, which includes the mass of protons, neutrons, and electrons (although electrons contribute very little to the total mass). Elements: 1. Definition: Elements are pure substances made up of only one type of atom. They are listed in the periodic table. 2. Properties: Each element has unique physical and chemical properties. These properties are determined by the arrangement of electrons in their atoms. 3. Symbol: Elements are represented by symbols, usually derived from their names. For example, hydrogen is represented by H, oxygen by O, and carbon by C. 4. Classification: Elements are classified into metals, nonmetals, and metalloids based on their properties and behaviour. Compounds: 1. Definition: Compounds are substances composed of two or more different elements chemically bonded together in fixed proportions. 2. Chemical Bonds: Compounds are formed through chemical bonding, which can be ionic, covalent, or metallic, depending on the types of elements involved. 3. Properties: Compounds have properties distinct from those of their constituent elements. For example, sodium (a metal) and chlorine (a nonmetal) combine to form sodium chloride (table salt), which is a compound with very different properties from either sodium or chlorine. 4. Naming: Compounds are named using specific rules depending on their type of bonding. Ionic compounds typically have names ending in "-ide," while covalent compounds use prefixes to indicate the number of atoms of each element. Common Names vs. Chemical Formulae: 1. Common Names: Many compounds have common names that are widely used in everyday language. These names are often based on historical or practical considerations. 2. Chemical Formulae: Chemical formulae are concise representations of compounds using symbols for the elements involved and numerical subscripts to indicate the ratio of atoms in the compound. Range of Compounds: 1. Inorganic Compounds: Water: Common Name - Water, Chemical Formula - H2O Table Salt: Common Name - Salt (Sodium Chloride), Chemical Formula - NaCl Baking Soda: Common Name - Sodium Bicarbonate, Chemical Formula - NaHCO3 Calcium Carbonate: Common Name - Calcium Carbonate (or chalk), Chemical Formula - CaCO3 Hydrochloric Acid: Common Name - Hydrochloric Acid, Chemical Formula - HCl Ammonia: Common Name - Ammonia, Chemical Formula - NH3 Sulfuric Acid: Common Name - Sulfuric Acid, Chemical Formula - H2SO4 2. Organic Compounds: Ethanol: Common Name - Alcohol (Ethanol), Chemical Formula - C2H5OH Glucose: Common Name - Glucose (Sugar), Chemical Formula - C6H12O6 Methane: Common Name - Methane, Chemical Formula - CH4 Acetic Acid: Common Name - Vinegar (Acetic Acid), Chemical Formula - CH3COOH Aspirin: Common Name - Aspirin (Acetylsalicylic Acid), Chemical Formula - C9H8O4 Sucrose: Common Name - Sugar (Sucrose), Chemical Formula - C12H22O11 Identifying Compounds: 1. Common Names: Recognize compounds by their common names, which are often used in everyday contexts. 2. Chemical Formulae: Understand the chemical composition of compounds through their chemical formulae, which provide information about the elements and their ratios in the compound. 3. Naming Rules: Learn the naming rules for different types of compounds (inorganic and organic) to identify compounds accurately based on their chemical formulas. 4. Functional Groups: In organic compounds, functional groups (such as hydroxyl groups in alcohols or carboxyl groups in acids) can help in identifying and naming compounds. Ionic Bonds: 1. Definition: Ionic bonds are formed between ions of opposite charges. One atom loses electrons to become a positively charged ion (cation), while another atom gains those electrons to become a negatively charged ion (anion). The electrostatic attraction between these ions holds them together in a compound.  2. Characteristics: Typically formed between a metal and a nonmetal. Strong electrostatic attraction between ions. Ionic compounds usually form crystalline structures and conduct electricity when dissolved in water or melted. 3. Examples of Ionic Compounds: Sodium Chloride (NaCl): Sodium (Na+) loses one electron to become a cation, while chlorine (Cl-) gains that electron to become an anion, resulting in the formation of NaCl. Calcium Oxide (CaO): Calcium (Ca2+) loses two electrons to form a cation, and oxygen (O2-) gains those electrons to form an anion, leading to the formation of CaO. Potassium Bromide (KBr): Potassium (K+) forms a cation by losing one electron, while bromine (Br-) forms an anion by gaining that electron, resulting in the compound KBr. Covalent Bonds: 1. Definition: Covalent bonds are formed when atoms share electrons to achieve a stable electron configuration. These bonds occur between nonmetal atoms. 2. Characteristics: Formed by sharing electrons between atoms. Can be polar (unequal sharing) or nonpolar (equal sharing) depending on the electronegativity of the atoms involved. Covalent compounds can exist as molecules with discrete units. 3. Examples of Covalent Compounds: Water (H2O): Oxygen shares electrons with hydrogen atoms to form covalent bonds, resulting in the molecule H2O. Carbon Dioxide (CO2): Carbon shares electrons with two oxygen atoms, forming covalent bonds and resulting in the molecule CO2. Methane (CH4): Carbon shares electrons with four hydrogen atoms, forming covalent bonds and resulting in the molecule CH4. Ammonia (NH3): Nitrogen shares electrons with three hydrogen atoms, forming covalent bonds and resulting in the molecule NH3. Differentiating Ionic and Covalent Bonds: Electron Sharing: Covalent bonds involve the sharing of electrons, while ionic bonds involve the transfer of electrons. Types of Atoms Involved: Covalent bonds typically occur between nonmetal atoms, while ionic bonds occur between a metal and a nonmetal. Bond Strength: Ionic bonds are generally stronger than covalent bonds due to the strong electrostatic attraction between ions.  Properties of Compounds: Ionic compounds often form crystalline structures and conduct electricity when dissolved in water, while covalent compounds can exist as discrete molecules and may not conduct electricity in their pure form. Types of Chemical Reactions: 1. Combustion Reactions: Definition: Combustion reactions are rapid chemical reactions that involve the reaction of a substance with oxygen, often producing heat and light. Examples: Combustion of hydrocarbons like methane (CH4) in the presence of oxygen to produce carbon dioxide (CO2) and water (H2O). Word Equation: Methane + Oxygen → Carbon Dioxide + Water Chemical Equation: CH4 + 2O2 → CO2 + 2H2O 2. Acid-Base Reactions: Definition: Acid-base reactions involve the transfer of protons (H⁺ ions) between acids and bases, resulting in the formation of water and a salt. Examples: Neutralization of hydrochloric acid (HCl) with sodium hydroxide (NaOH) to produce water (H2O) and sodium chloride (NaCl). Word Equation: Hydrochloric Acid + Sodium Hydroxide → Water + Sodium Chloride Chemical Equation: HCl + NaOH → H2O + NaCl 3. Redox (Oxidation-Reduction) Reactions: Definition: Redox reactions involve the transfer of electrons between reactants. Oxidation involves the loss of electrons, while reduction involves the gain of electrons. Examples: Rusting of iron (Fe) in the presence of oxygen (O2), where iron loses electrons (oxidation) and oxygen gains electrons (reduction). Word Equation: Iron + Oxygen → Iron Oxide Chemical Equation: 4Fe + 3O2 → 2Fe2O3 4. Synthesis (Combination) Reactions: Definition: Synthesis reactions involve the combination of two or more substances to form a single compound. Examples: Formation of water (H2O) from hydrogen gas (H2) and oxygen gas (O2). Word Equation: Hydrogen + Oxygen → Water Chemical Equation: 2H2 + O2 → 2H2O 5. Decomposition Reactions: Definition: Decomposition reactions involve the breakdown of a compound into simpler substances. Examples: Decomposition of hydrogen peroxide (H2O2) into water (H2O) and oxygen gas (O2). Word Equation: Hydrogen Peroxide → Water + Oxygen Chemical Equation: 2H2O2 → 2H2O + O2 6. Photosynthesis (Living System):  Definition: Photosynthesis is a complex chemical process in plants and some microorganisms where light energy is used to convert carbon dioxide (CO2) and water (H2O) into glucose (C6H12O6) and oxygen gas (O2). Word Equation: Carbon Dioxide + Water + Light Energy → Glucose + Oxygen Chemical Equation: 6CO2 + 6H2O + Light Energy → C6H12O6 + 6O2 7. Cellular Respiration (Living System): Definition: Cellular respiration is a biochemical process in living organisms where glucose (C6H12O6) is oxidised to release energy, carbon dioxide (CO2), and water (H2O). Word Equation: Glucose + Oxygen → Carbon Dioxide + Water + Energy Chemical Equation: C6H12O6 + 6O2 → 6CO2 + 6H2O + Energy Acids and Bases: 1. Acids: Definition: Acids are substances that can donate protons (H⁺ ions) in aqueous solutions. They have a pH less than 7 and turn blue litmus paper red. Chemical Formulae: Common acids include hydrochloric acid (HCl), sulfuric acid (H2SO4), nitric acid (HNO3), and acetic acid (CH3COOH). 2. Bases: Definition: Bases are substances that can accept protons or donate hydroxide ions (OH⁻) in aqueous solutions. They have a pH greater than 7 and turn red litmus paper blue. Chemical Formulae: Common bases include sodium hydroxide (NaOH), potassium hydroxide (KOH), calcium hydroxide (Ca(OH)2), and ammonia (NH3). Reactions Involving Acids and Bases: 1. Acid-Base Neutralization: Definition: Acid-base neutralisation reactions occur when an acid reacts with a base to form water and a salt. Word Equation: Acid + Base → Water + Salt Chemical Equation: HCl + NaOH → H2O + NaCl 2. Formation of Salts: Definition: Salts are compounds formed from the reaction between an acid and a base, where the hydrogen ions (H⁺) from the acid are replaced by metal or ammonium ions. Word Equation: Acid + Base → Salt + Water Chemical Equation: H2SO4 + 2KOH → K2SO4 + 2H2O 3. Acid-Base Reactions with Metals: Definition: Acids react with metals to form salts and hydrogen gas (H2). Word Equation: Acid + Metal → Salt + Hydrogen Gas Chemical Equation: 2HCl + Zn → ZnCl2 + H2 4. Neutralisation of Carbonates and Hydrogen Carbonates:  Definition: Acids neutralise carbonates and hydrogen carbonates (bicarbonates) to form carbon dioxide gas (CO2), water (H2O), and salt. Word Equation: Acid + Carbonate/Bicarbonate → Carbon Dioxide + Water + Salt Chemical Equation: 2HCl + Na2CO3 → 2NaCl + H2O + CO2 5. Reaction with Ammonia: Definition: Acids react with ammonia (a weak base) to form ammonium salts. Word Equation: Acid + Ammonia → Ammonium Salt Chemical Equation: HCl + NH3 → NH4Cl Importance of Acid-Base Reactions: Industrial Applications: Acid-base reactions are used in various industries such as chemical manufacturing, water treatment, and food processing. Environmental Impact: Understanding acid-base reactions is crucial for addressing environmental issues like acid rain. Biological Functions: Acid-base balance is essential for physiological processes in living organisms, including digestion and cellular functions. Law of Conservation of Mass: 1. Definition: The Law of Conservation of Mass states that in a closed system, mass is neither created nor destroyed during a chemical reaction. This means that the total mass of the reactants is equal to the total mass of the products. 2. Explanation: According to this law, atoms are neither created nor destroyed during a chemical reaction. Instead, they are rearranged to form new compounds. 3. Example: In the reaction between hydrogen gas (H2) and oxygen gas (O2) to form water (H2O), the total mass of hydrogen and oxygen in the reactants is equal to the total mass of hydrogen and oxygen in the products. Number of Atoms in Chemical Reactions: 1. Balancing Chemical Equations: Balancing chemical equations is the process of ensuring that the number of atoms of each element is the same on both sides of the equation. This is necessary to obey the Law of Conservation of Mass. 2. Example: The combustion of methane (CH4) with oxygen (O2) to form carbon dioxide (CO2) and water (H2O) can be represented as: Unbalanced Equation: CH4 + O2 → CO2 + H2O Balanced Equation: CH4 + 2O2 → CO2 + 2H2O In this balanced equation, there are 1 carbon atom, 4 hydrogen atoms, and 4 oxygen atoms on both sides, preserving mass. 3. Coefficients in Equations: Coefficients in chemical equations represent the number of molecules or moles of each substance involved in the reaction. These coefficients are used to balance equations and ensure conservation of mass.  4. Conservation of Atoms: The Law of Conservation of Mass also implies the conservation of atoms. This means that the total number of each type of atom on the reactant side must equal the total number of that atom on the product side. Importance of Conservation of Mass: 1. Predicting Products: Understanding the conservation of mass helps predict the products of chemical reactions based on the number of atoms of each element involved. 2. Experimental Verification: The Law of Conservation of Mass has been experimentally verified countless times and is a fundamental principle in chemistry. 3. Stoichiometry Calculations: Conservation of mass is essential in stoichiometry calculations, where the quantities of reactants and products in a chemical reaction are determined based on the balanced equation. Part C Basic Structure of Atoms: 1. Nucleus: The nucleus is the central core of an atom. It contains protons and neutrons. Protons have a positive charge, while neutrons are neutral. 2. Electron Cloud: Surrounding the nucleus is the electron cloud. Electrons are negatively charged particles found in the electron cloud. Electrons orbit the nucleus in energy levels or shells. Subatomic Particles: 1. Protons: Positively charged particles found in the nucleus. The number of protons determines the element's identity (atomic number). The atomic number is represented by "Z" in the periodic table. 2. Neutrons: Neutral particles found in the nucleus. They help stabilise the nucleus and contribute to the atomic mass. The sum of protons and neutrons gives the mass number of an atom. 3. Electrons: Negatively charged particles found in the electron cloud. They occupy specific energy levels or shells around the nucleus. The number of electrons is equal to the number of protons in a neutral atom. Energy Levels or Shells: 1. First Energy Level (K Shell): Closest to the nucleus. Can hold up to 2 electrons. 2. Second Energy Level (L Shell): Located further from the nucleus. Can hold up to 8 electrons. 3. Third Energy Level (M Shell): Further from the nucleus than the second energy level. Can hold up to 18 electrons. Electron Configuration: 1. Definition: Electron configuration refers to the arrangement of electrons in the energy levels or shells of an atom. It follows the Aufbau principle, Pauli exclusion principle, and Hund's rule. 2. Example: Sodium (Na) has an atomic number of 11. Its electron configuration is 2-8-1, meaning 2 electrons in the first energy level, 8 in the second, and 1 in the third. Isotopes: 1. Definition: Isotopes are atoms of the same element with different numbers of neutrons. They have the same number of protons (same atomic number) but different mass numbers. 2. Example: Carbon has isotopes like Carbon-12 (6 protons and 6 neutrons) and Carbon-14 (6 protons and 8 neutrons). Importance of Atom Structure: 1. Chemical Properties: The structure of atoms determines their chemical properties and reactivity. Elements with similar electron configurations exhibit similar chemical behaviours. 2. Nuclear Reactions: Understanding atomic structure is crucial in nuclear reactions involving changes in the nucleus (e.g., nuclear fusion, fission). 3. Electronic Configuration: Electron configuration influences bonding patterns, molecular shapes, and properties of compounds. 1. Dalton's Model (Solid Sphere Model): Proposed by: John Dalton (early 19th century). Description: Atoms were considered as indivisible and solid spheres, like billiard balls. Key Points: All atoms of a given element are identical in mass and properties. Atoms of different elements have different masses and properties. Atoms combine in simple whole number ratios to form compounds. 2. J.J. Thomson's Model (Plum Pudding Model): Proposed by: J.J. Thomson (late 19th century). Description: Atoms were viewed as a positively charged sphere with negatively charged electrons embedded in it, like raisins in a pudding. Key Points: Introduced the concept of electrons as subatomic particles. Proposed a neutral overall charge for the atom. 3. Rutherford's Model (Planetary Model): Proposed by: Ernest Rutherford (early 20th century). Description: Atoms have a central nucleus containing protons and neutrons, with electrons orbiting around the nucleus similar to planets orbiting the sun. Key Points: Discovered the nucleus through the gold foil experiment. Majority of the atom's mass and positive charge is concentrated in the nucleus. Electrons orbit the nucleus in fixed paths or orbits. 4. Bohr's Model (Bohr-Rutherford Model): Proposed by: Niels Bohr (early 20th century). Description: Electrons move in specific energy levels or shells around the nucleus, and each shell has a fixed energy level. Key Points: Electrons can move between energy levels by absorbing or emitting energy. Only certain orbits are allowed, called quantized orbits. Explained atomic spectra and the stability of atoms. 5. Quantum Mechanical Model (Modern Model): Developed by: Various scientists, including Schrödinger, Heisenberg, and others (early to mid-20th century). Description: Describes electrons as having wave-like properties and is based on the principles of quantum mechanics. Key Points: Electrons are described by wave functions or orbitals, which represent the probability of finding an electron in a specific region around the nucleus. Introduced the concept of electron clouds or probability distributions rather than fixed orbits. Provides a more accurate description of atomic structure and behaviour. Importance of Different Models: 1. Evolution of Understanding: Each model represents an advancement in our understanding of the atom, building upon the previous model's successes and addressing its limitations. 2. Technological Advances: Advances in experimental techniques and theoretical frameworks led to the development of more sophisticated models. 3. Applications: Different models have practical applications in fields like chemistry, physics, and engineering, influencing research, technology, and material design. Radioactivity and Atoms Radioactivity: Definition: Radioactivity is the spontaneous emission of particles or radiation from the unstable nucleus of an atom. Causes of radioactivity: Unstable atomic nuclei: Radioactive elements have unstable nuclei that undergo decay to achieve a more stable configuration. Nuclear reactions: Certain nuclear reactions, such as fission and fusion, can also produce radioactive isotopes. Radioactive Elements and Isotopes: Radioactive elements: These are elements that have isotopes with unstable nuclei, leading to the emission of radiation. Isotopes: Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons. Some isotopes of elements can be stable, while others are radioactive. Half-life of Radioactive Isotopes: Definition: The half-life of a radioactive isotope is the time it takes for half of the nuclei in a sample of the isotope to undergo radioactive decay. Characteristics: Different isotopes have different half-lives, ranging from fractions of a second to billions of years. Half-life is a characteristic property of each radioactive isotope and is not affected by external conditions such as temperature or pressure. The concept of half-life is used in radiometric dating to determine the age of rocks and fossils based on the decay of radioactive isotopes within them. Isotopes in Industry Carbon-14 (14C): Use: Carbon-14 is used in radiocarbon dating to determine the age of archaeological artefacts and geological samples. Explanation: This isotope decays at a known rate, allowing scientists to estimate the age of organic materials based on the remaining amount of Carbon-14. Cobalt-60 (60Co): Use: Cobalt-60 is used in radiotherapy to treat cancerous tumours. Explanation: Its high energy gamma rays can destroy cancer cells, making it an effective tool in oncology. Uranium-235 (235U): Use: Uranium-235 is used as fuel in nuclear reactors to generate electricity. Explanation: Its fission properties release large amounts of energy, which is harnessed to produce electricity through nuclear reactions. Advantages and Disadvantages of Using Isotopes in Industries: Advantages: Precision: Isotopes offer precise measurements and calculations, crucial in scientific research and medical diagnostics. Efficiency: In industries like energy production, isotopes can generate large amounts of energy efficiently, reducing reliance on traditional fossil fuels. Innovation: Isotope applications drive innovation, leading to advancements in medicine, agriculture, and environmental studies. Disadvantages: Safety Concerns: Some isotopes, if mishandled, can pose radiation hazards, requiring strict safety protocols. Cost: Isotope production and handling can be expensive, limiting widespread adoption in some industries. Waste Management: Disposal of radioactive isotopes generates nuclear waste, necessitating safe storage and disposal methods. Applications of Half-Life and Carbon - 14 Half-Life Basics: Half-life refers to the time required for half of a quantity of a substance to undergo decay or transformation. It's a fundamental concept in nuclear physics and is used to describe the stability or decay rate of radioactive substances. Carbon-14 Dating: Carbon-14 (C-14) dating is a prominent application of half-life in archaeology and geology. C-14 is a radioactive isotope of carbon that decays over time. It's produced in the upper atmosphere when cosmic rays interact with nitrogen atoms. Living organisms continuously absorb carbon, including C-14, from the atmosphere or food sources. When an organism dies, it stops absorbing C-14, and the existing C-14 begins to decay. By measuring the remaining C-14 in organic samples like bones or wood, scientists can estimate the age of the sample since death, providing valuable information for dating archaeological artefacts or determining geological ages. Limitations of Carbon-14 Dating: Carbon-14 dating is effective for samples up to around 50,000 years old. Beyond this timeframe, the amount of remaining C-14 becomes too low to measure accurately. It's not suitable for dating materials that are not organic, like rocks or metals, as they don't absorb atmospheric carbon during their formation. Medical Applications: Half-life concepts are vital in medical imaging techniques like positron emission tomography (PET) scans. Radioactive isotopes with short half-lives are used as tracers in PET scans to visualise metabolic processes and detect abnormalities in organs or tissues. Industrial and Environmental Uses: Half-life calculations are crucial in managing radioactive waste from nuclear power plants. Understanding the decay rates of radioactive isotopes helps in determining safe storage times and methods for radioactive waste. Environmental scientists also use half-life principles to study the impact of radioactive substances on ecosystems and human health. Archaeological Insights: Beyond carbon dating, half-life principles are applied in other dating methods like potassium-argon dating and uranium-lead dating, used for older geological samples. These methods rely on measuring the decay of isotopes with longer half-lives, providing insights into Earth's geological history and the evolution of life on the planet. Future Developments:  Ongoing research aims to refine dating techniques, extending their applicability to older samples or developing new methods for non-organic materials. Advances in technology, such as accelerator mass spectrometry, have improved the accuracy and sensitivity of carbon dating, opening new avenues for archaeological and environmental studies Living with Radiation Outcome 1: Students can identify the use of Radiation Key Concepts: Radiation Types: ○ Ionizing Radiation: Includes alpha particles, beta particles, gamma rays, and X-rays. ○ Non-Ionizing Radiation: Includes ultraviolet (UV) light, visible light, infrared radiation, microwaves, and radio waves. Common Uses: ○ Medical Applications: Diagnostic Imaging: X-rays, CT scans, and PET scans. Radiation Therapy: Treatment for cancer and other diseases. ○ Industrial Applications: Non-destructive Testing: Inspecting welds and materials for structural integrity. Radiography: Checking for faults in materials. ○ Energy Production: Nuclear Power Plants: Using nuclear fission to generate electricity. ○ Consumer Products: Smoke Detectors: Use of americium-241. Luminous Watches: Use of tritium or radium. ○ Scientific Research: Radiotracers: Used in biological and chemical research. Outcome 2: Students can research and analyse the effects of using radiation in everyday life Key Concepts: Biological Effects: ○ Short-term Exposure: Effects like radiation burns, acute radiation syndrome. ○ Long-term Exposure: Increased risk of cancer, genetic mutations, cataracts. Environmental Impact: ○ Radioactive Contamination: Impact on air, water, and soil. ○ Nuclear Accidents: Historical examples like Chernobyl and Fukushima. Safety and Regulation: ○ Regulatory Bodies: Roles of organisations like the International Atomic Energy Agency (IAEA) and the Environmental Protection Agency (EPA). ○ Safety Standards: Guidelines for safe levels of radiation exposure.  Societal Impact: ○ Medical Advancements: Improved diagnostic and treatment methods. ○ Public Health Concerns: Balancing benefits and risks of radiation use. ○ Ethical Considerations: Debates over nuclear energy and medical uses. Uses in radioisotopes industry Radi Which isotope it is and Its half life Type of Applications and Interesting Disadvantages of oacti what it changes to radiation reasons for using facts using isotopes ve emitted the isotope Isoto pes Co Naturally occurring Co is a Gamma Co-60 is used Cobalt Because it decays cobalt (27Co) consists nuclear Rays medically for (chemical by gamma of a single stable isomer of radiation therapy symbol Co) is radiation, external isotope, 59Co (thus, 60Co with a as implants and as a hard, exposure to large cobalt is a mononuclidic half-life of an external source grey-blue sources of Co-60 element). Twenty-eight 10.467 of radiation metal that is can cause skin radioisotopes have minutes. It exposure. It is solid under burns, acute been characterised; the decays by used industrially in normal radiation sickness, most stable are 60Co internal levelling gauges conditions. or death. Most with a half-life of 5.2714 years, 57Co (271.8 transition and to x-ray Cobalt is Co-60 that is days), 56Co (77.27 to 60Co, welding seams and similar to ingested is days), and 58Co (70.86 emitting other structural iron and excreted in the days). 58.6 keV elements to detect nickel in its faeces; however, a gamma flaws. properties small amount is rays, or and can be absorbed by the with a low magnetised liver, kidneys, and probability like iron. The bones. (0.22%) by most β-decay common into 60Ni. radioactive isotope. For example, uranium has thirty-seven different isotopes, including uranium-235 and uranium-238. I Iodine-131 is an important It has a The most It is associated I-131 is used Exposure to I-131 radioisotope of iodine radioactive widely with nuclear in medicine can increase the discovered by Glenn decay used energy, medical to diagnose risk of thyroid Seaborg and John half-life of iodine diagnostic and and treat cancer. It is Livingood in 1938 at the radioisot treatment cancers of thought that risk is  University of California, about eight ope, procedures, and the thyroid higher for people Berkeley. days. iodine-1 natural gas gland. Where who have had 31, emits production. does it come multiple exposures radiation from? I-131 is and for people in the produced exposed at a form of commercially younger age. But medium for medical even among energy and industrial people who have gamma uses through documented rays and nuclear exposures to I-131, beta fission. It also few develop this particles, is a cancer. which byproduct of disrupts nuclear molecule fission s in cells processes in and nuclear deposits reactors and energy in weapons tissues, testing. causing damage. Na There are thirteen Periodic Gamma Na is a There is only Sodium has recognized isotopes of Table--Sodi Rays positron-emitting one stable drawbacks sodium. Na is the only um. Sodium isotope with a isotope of associated with its stable isotope. As such, it has two remarkably long sodium: intense chemical is considered a radioactive half-life. It is used 23Na. The activity (catching monoisotopic element cosmogenic to create symbol for fire) in the case of and it has a standard isotopes test-objects and sodium is Na, a leak to the atomic mass: (22Na, point-sources for which comes atmosphere. This 22.98976928(2). half-life = positron emission from the leads to the 2.605 tomography. Latin natrium incorporation of a years; or Arabic or a double envelope 24Na, similar-sound for the primary half-life = ~ ing Egyptian circuit with an 15 hours) word, all inert interspace. that have referring to been used soda or as tracers in sodium hydrologic carbonate. studies. Sodium is an abundant element. It's found in the sun and many other stars. Sr Periodic Table--Strontium. The half-life Sr-90 Since different The Because of the The alkali earth metal of decays to geological strontium beta radiation strontium has four stable, strontium-9 yttrium formations present isotope Sr-90 emitted from the naturally occurring 0 (the time 90 specific Sr isotopic is a nuclear Isotope Sr, Y-90 isotopes: 84Sr (0.56%), it takes for (Y-90), signatures, the fission poses a risk of 86Sr (9.86%), 81Sr (7.0%) half of the which in 87Sr/86Sr ratio product and burns to the eyes and 88Sr (82.58%). The strontium turn can be used to is released and on the skin isotopes of Sr include = to give off decays study oil migration into the from external Strontium-90, its radiation by beta and mixing environment exposure. Strontium-87, and change radiation between during Strontium-86, into so that geological units nuclear Strontium-88, another whereve and their fallout. Sr-90 Strontium-80, substance) r Sr-90 is compartmentalizat has a half-life strontium-85, is 29 years. present ion efficiency. The of around 28 Strontium-95, Y-90 is Sr isotopic ratio is years. When Strontium-96, also also broadly used high levels of Strontium-89. present. for production Sr-90 are water tracing and absorbed by isotope bone tissue stratigraphy. in place of calcium, it can destroy bone marrow and cause cancer. U Uranium is a chemical Uranium is All U-235 A uranium Inhaling large element with the symbol a isotopes concentrations can atom has 92 concentrations of U and the atomic number radionuclid of be used as fuel for protons and uranium can cause 92. There are three e that has uranium nuclear power 92 electrons, lung cancer from naturally occurring an are plants and the of which 6 the exposure to isotopes of uranium: extremely radioacti nuclear reactors are valence alpha particles. uranium-238, the heaviest long ve. Both that run naval electrons. Uranium is also a and most abundant, half-life. uranium ships and Uranium has toxic chemical, uranium-235 and Naturally and submarines. It also the highest meaning that uranium-234. occurring depleted can be used in atomic ingestion of uranium-23 uranium, nuclear weapons. weight of all uranium can cause 8 present in and their Whilst naturally kidney damage the Earth's immedia Uranium-238 is occurring from its chemical crust has a te decay majorly used for elements. properties much half-life of products Paleontology and Uranium sooner than its almost 4.5 , emit radiometric dating occurs radioactive billion alpha in fossils. naturally in properties would years. If you and beta low cause cancers of take a soil particles concentratio the bone or liver. sample and a ns in soil, anywhere small rock and in the amount water, and is world, of commercially  including gamma extracted your radiation from backyard,. uranium-bear you will ing minerals find such as uranium uraninite. atoms that date back to when the Earth was formed. Am Americium is produced The time in Am-241 Most americium is Americium is Americium poses a when plutonium absorbs which half is produced by a shiny silver more significant neutrons in nuclear the atoms primarily uranium or radioactive risk if ingested reactors or during nuclear of a an alpha plutonium being metal. All (swallowed) or weapons tests. radioactive emitter, bombarded with isotopes of inhaled. Once in Americium-241 is the substance but also neutrons in this element the body, it tends most common form of disintegrate emits nuclear reactors – are to concentrate in Americium. Additionally to another some one tonne of spent radioactive. the bone, liver, and Americium has = nuclear gamma nuclear fuel The isotope muscle. Americium Americium-242m, form is rays. contains about 100 with the can stay in the Americium-241, known as grams of longest body for decades Americium-239 and the half-life. americium. It is half-life is and continue to Americium-243. The half-life widely used in americium-2 expose the of commercial 43, which has surrounding americium- ionisation a half-life of tissues to 241 is chamber smoke 7370 years. radiation, about 432 detectors, as well The most increasing the risk years. as in neutron common of developing sources and isotopes are cancer. industrial gauges. americium-2 41, with a half-life of 432.7 years, and americium-2 43. Pu Plutonium is a chemical Pu-238 with Because Plutonium-238 Plutonium is Poor thermal element with symbol Pu a half-life of it emits generates named for capability. Poor and atomic number 94. 88 years alpha significant heat the dwarf weatherability. Classified as an actinide, has a particles, through its planet Pluto. Attacked by most Plutonium is a solid at relatively plutoniu radioactive decay Plutonium is solvents. Utilise room temperature. high heat m is process, which not a good toxic isocyanates. Pu can change to: production most makes it useful as conductor of Flammable. Plutonium-239 rate which dangero a heat source for electricity or Plutonium-240 makes it us when sensitive electrical heat, unlike Plutonium-238 useful as a inhaled. components in some metals.  Plutonium-241 power When satellites, as well The alpha Uranium-235 source with plutoniu as a power source form of Uranium-238 a long m (for example, plutonium is Plutonium-242 service life. particles battery power) for hard and Plutonium-243 are satellites. brittle, while Plutonium-236 inhaled, Plutonium-239 is the delta Plutonium-231 they used to make form is soft Plutonium-229 lodge in nuclear weapons. and ductile. Plutonium-234 the lung Plutonium Plutonium-244 tissue. occurs Uranium-240 The naturally in Isotopes of neptunium alpha the Earth's Neptunium-238 particles crust in Plutonium-232 can kill uranium Plutonium-237 lung ores, but it is Plutonium-245 cells, very rare. Plutonium-233 which Plutonium-230 causes Uranium-239 scarring Plutonium-246 of the Plutonium-235 lungs, Plutonium-228 leading Plutonium-247 to further lung disease and cancer. Tc All the isotopes of 6.04 hours. Gamma- Technetium is a It is the Most commonly, technetium are rays remarkable lightest technetium-99m radioactive. It is one of corrosion inhibitor element causes rash, two elements with Z < 83 for steel, and whose angioedema, fever, that have no stable adding very small isotopes are and anaphylaxis isotopes; the other amounts can all due to element is promethium (Z provide excellent radioactive. hypersensitivity = 61). Technetium has protection. This Technetium reactions. Patients three long-lived use is limited to and may also radioactive isotopes: 97Tc closed systems as promethium experience a (T1/2 =2.6 x 106 years), technetium is are the only transient increase 98Tc (T1/2 = 4.2 x 106 radioactive. radioactive in blood pressure, years) and 99Tc (T1/2 = Technetium has no elements seizures, 2.1 x 105 years). Tc can known biological whose arrhythmias, and change to: role. It is toxic due neighbours in syncope. When Technetium-99 to its radioactivity. the sense of used in abdominal Technetium-99m atomic imaging, Technetium-98 number are abdominal pain, Technetium-97 both stable. vomiting, and Technetium-101  Technetium-100 diarrhoea may Technetium-92 occur. Atomic Theory Date Scientist Contribution to Atomic Theory Picture Democritus Democritus was a Greek philosopher (470-380 400 BC B.C.) who is the father of modern atomic thought. He proposed that matter could NOT be divided into smaller pieces forever. He claimed that matter was made of small, hard particles that he called “atomos” John Dalton created the very first atomic 1807 John Dalton theory. Dalton viewed atoms as tiny, solid balls. Dalton was an English school teacher who performed many experiments on atoms. His atomic theory had 4 statements: 1. Atoms are tiny, invisible particles. 2. Atoms of one element are all the same. 3. Atoms of different elements are different. 4. Compounds form by combining atoms. J.J. Thomson discovered electrons. He also 1903 JJ Thompson proposed the existence of a (+) particle. His atomic model was known as the “raisin bun model”.  Ernest Rutherford conducted a famous 1911 Ernest experiment called the gold foil experiment. He Rutherford used a thin sheet of gold foil. He also used special equipment to shoot alpha particles (positively charged particles) at the gold foil. Most particles passed straight through the foil like the foil was not there. Some particles went straight back or were deflected (went in another direction) as if they had hit something. The experiment shows: Atoms are made of a small positive nucleus; positive nucleus repels (pushes away) positive alpha particles, Atoms are mostly empty space Niels Bohr was a Danish physicist. He 1913 Niels Bohr proposed a model of the atom that is similar to the model of the solar system. The electrons go around the nucleus like planets orbit around the sun. All electrons have their energy levels – a certain distance from the nucleus. Each energy level can hold a certain number of electrons. Level 1 can hold 2 electrons, Level 2 - 8 electrons, Level 3 - 18 electrons, and level 4 – 32 electrons. The energy of electrons goes up from level 1 to other levels. When electrons release (lose) energy they go down a level. When electrons absorb (gain) energy, they go to a higher level. James Chadwick's contribution to the atomic 1932 James model was his discovery of the neutron. The Chadwick neutron is a neutrally charged subatomic particle that is about the same mass as the proton. Both protons and neutrons occupy the nucleus of the atom. Chadwick was able to discover the neutron and measure its mass. Coordination and Disease Identify the five senses and its importance. There are five basic human senses: touch, sight, hearing, smell and taste. The sensing organs associated with each sense send information to the brain to help us understand and perceive the world around us. Define the three adaptations. The three types of adaptation include structural, physiological, and behavioural. Structural adaptation results in a change in physical appearance. Physiological adaptation results in biological changes on a cellular level. Behaviour adaptations result from adapted behaviour based on environmental stimuli. Identify the building blocks of the living hierarchy. A cell is the smallest unit of a living thing. A living thing, whether made of one cell (like bacteria) or many cells (like a human), is called an organism. Thus, cells are the basic building blocks of all organisms. Identify and describe two examples of the organ systems in the human body. 1) Lymphatic- Defends against infection and disease and transfers lymph between tissues and the blood stream. 2) Digestive - Processes foods and absorbs nutrients, minerals, vitamins, and water. Pictures: Cell: Cells are the basic building blocks of all living things. The human body is composed of trillions of cells. They provide structure for the body, take in nutrients from food, convert those nutrients into energy, and carry out specialised functions.  : Tissue: Your body is made of cells and when groups of cells do the same kind of work, they are called tissues. You have four main types of tissues: Connective, Epithelial, Muscle, and Nervous tissue. Connective tissue joins bones and cushions organs. Epithelial tissue covers the outside of the body. : Organ: Organs are the body's recognizable structures (for example, the heart, lungs, liver, eyes, and stomach) that perform specific functions. An organ is made of several types of tissue and therefore several types of cells. Human Organ Systems: An organ system is a group of organs that work together to perform one or more functions. Each organ has a specific function in the body and is made up of certain tissues. 1. Below is the human body. In the next page choose three organ systems and describe its functions. Also label the organ systems using the Labels on the body. Use the shapes and text box function to answer the following questions if done on the computer. 1) Brain: The brain is a complex organ that controls thought, memory, emotion, touch, motor skills, vision, breathing, temperature, hunger and every process that  regulates our body. Together, the brain and spinal cord that extends from it make up the central nervous system, or CNS. 2) Lungs: The lungs and respiratory system allow us to breathe. They bring oxygen into our bodies (called inspiration, or inhalation) and send carbon dioxide out (called expiration, or exhalation). This exchange of oxygen and carbon dioxide is called respiration. 3) Heart: The functions of the heart are to pump blood and oxygen around the body and deliver waste products (carbon dioxide) back to the lungs to be removed. The heart consists of four chambers, each separated by valves which direct the flow of blood. 4) Intestines: Its main purpose is to digest food. But the intestine is not only there for digestion: It also produces various hormones that carry messages to other parts of the body, and plays an important role in fighting germs and regulating water entering and leaving the body. 5) Muscles: Your leg muscles help you move, carry the weight of your body and support you when you stand. You have several muscles in your upper and lower legs. They work together to enable you to walk, run, jump and flex and point your feet. Table below: Organ System: Description of its Functions: Brain The brain is a complex organ that controls thought, memory, emotion, touch, motor skills, vision, breathing, temperature, hunger and every process that regulates our body. Together, the brain and spinal cord that extends from it make up the central nervous system, or CNS. Lungs The lungs and respiratory system allow us to breathe. They bring oxygen into our bodies (called inspiration, or inhalation) and send carbon dioxide out (called expiration, or exhalation). This exchange of oxygen and carbon dioxide is called respiration. Heart The functions of the heart are to pump blood and oxygen around the body and deliver waste products (carbon dioxide) back to the lungs to be removed. The heart consists of four chambers, each separated by valves which direct the flow of blood. Intestines Its main purpose is to digest food. But the intestine is not only there for digestion: It also produces various hormones that carry messages to other parts of the body, and plays an important role in fighting germs and regulating water entering and leaving the body. Muscles Your leg muscles help you move, carry the weight of your body and support you when you stand. You have several muscles in your upper and lower legs. They work together to enable you to walk, run, jump and flex and point your feet. Explain what homeostasis is and outline why it is important to survival: Homeostasis is any self-regulating process by which an organism tends to maintain stability while adjusting to conditions that are best for its survival. If homeostasis is successful, life continues; if it's unsuccessful, it results in a disaster or death of the organism. Identify a chemical and physical change an organism can react to: In a chemical reaction, there is a change in the composition of the substances in question; in a physical change there is a difference in the appearance, smell, or simple display of a sample of matter without a change in composition. Although we call them physical "reactions," no reaction is actually occurring. Organisms for chemical change can react to changes in composition of their foods and for physical changes they can react to a changing environment. The Brain Nervous System Endocrine System - Composed of the Brain, Spinal Cord and Nerves - Composed of glands - Acts via nervous impulses travelling along the - Acts via the release of hormones into the neurons bloodstream - Requires diffusion of neurotransmitters - Requires diffusion of hormones from the blood into target cells and tissues  - Provides relatively faster responses (reflex arcs - Provides relatively longer lasting responses being the fastest) - Causes effectors to respond e.g. muscles to - Causes targeted responses, e.g. growth, control of contract, or glands to secrete hormones metabolic reactions, development and reproduction Additional table: Nervous System Endocrine System - Composed of: The nervous system is - Composed of: The endocrine system made up of the central nervous includes the hypothalamus, pineal system and the peripheral nervous gland, pituitary gland, thyroid gland, system: The central nervous system parathyroid glands, thymus, adrenal includes the brain and spinal cord. glands, and pancreas. It also includes The peripheral nervous system the testes in males and the ovaries includes the nerves that run and placenta (during pregnancy) in throughout the whole body. females. - Acts Via: The nervous system uses - Acts Via: Endocrine glands make tiny cells called neurons to send chemicals called hormones and pass messages back and forth from the them straight into the bloodstream. brain, through the spinal cord, to the Hormones can be thought of as nerves throughout the body. Billions of chemical messages. From the neurons work together to create a bloodstream, the hormones communication network. communicate with the body by heading towards their target cell to bring about a particular change or effect to that cell. - Diffusion: Requires diffusion in the - Diffusion: Endocrine communication formation of morphogen gradients at involves chemical signalling via the the time of embryogenesis, and for its release of hormones into the basic function, e.g. in the transfer of extracellular fluid. From there, chemical signals from one neuron to hormones diffuse into the another during neurotransmission. bloodstream and may travel to distant body regions, where they elicit a response in target cells. Endocrine glands are ductless glands that secrete hormones. - Provides: The nervous system is - Provides: Endocrine glands release made up of the brain, spinal cord and hormones into the bloodstream. This nerves. It controls many aspects of lets the hormones travel to cells in  what you think, how you feel and what other parts of the body. The endocrine your body does. It allows you to do hormones help control mood, growth things such as walk, speak, swallow, and development, the way our organs breathe and learn. It also controls how work, metabolism , and reproduction. the body reacts in stressful situations. - Causes: Structural disorders, such as - Causes: The endocrine system affects brain or spinal cord injury, Bell's palsy, how your heart beats, how your bones cervical spondylosis, carpal tunnel and tissues grow, and even your ability syndrome, brain or spinal cord to make a baby. Disorders of the tumours, peripheral neuropathy, and endocrine system happen if your Guillain-Barré syndrome. Functional hormone levels are too high or too disorders, such as headache, epilepsy, low, or if your body doesn't respond to dizziness, and neuralgia. hormones in the expected way. Part