Honors Bio: Biodiversity and Life
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Biology = the observation, exploration, and investigation of life
Although distinguishing between living and not is often an innate ability, actually defining what life means involves many different explanations and statements about what an organism does, and the actions they take.
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Cell = the first unit of life that all living things are composed of
Organization: every organism must have hierarchical levels of their structures
Reproduction: ability to reproduce
Growth and Development: hereditary instructions that dictate growth
Response to the environment: ability to react to external stimuli
Energy Processing: intake and utilization of energy
Regulation: innate mechanisms that maintain internal balances
All other organisms’ properties of life originate from cells, where maintenance of internal balances, processing energy in and out, sensing and reacting to stimuli, and the ability to reproduce (each cell rises from another cell–important not only for reproduction but for healing injured tissue as well) form the basis all life performs. Organisms can be uni or multi-celled, and each must meet these requirements of life.
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Biosphere = inclusion of all life and their environments
Ecosystem = a certain area’s life presence and their corresponding habitats/environmental interactions
Community = all of an ecosystem’s present life
Population = every member of a specific species within a specific area
Biologists study the many levels of life, which organize it in a structured and understandable way.
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Emergent properties = new and previously absent properties which only appear at the next level
Ex:
The neural signals transmitted through an organism’s body is a product of how its nervous system is arranged with every other organ system within the body
Organism = one single living being
Organ/organ system = areas or components of body with specific functions; similarly functioning organs may work together to form an organ system
Tissue = similarly functioning groups of cells
Cell = smallest unit of life (most foundational unit)
Organelle = specialized, membrane-enclosed parts of cell
Molecule = two or more atoms bonded together
Beginning from molecules and their basic chemical principles, new properties of life arise from each new level of organization. This is a result of the accumulation and culmination of each new level. Emergent properties represent these additional and crucial characteristics.
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Matter = substance that takes up space and has volume
Everything material in the universe is composed of matter, and exists as either liquid, solid, or gas.
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Element = a substance that, through non-chemical means, cannot be broken down any further
How matter exists in the world has an incredibly wide range. An element is used to describe substances in their most basic forms, which means that nothing chemical can break them down any further. With 92 naturally occurring elements, some scientists have managed to create a few synthetic elements, which are labeled with the first two letters of their German, Latin, or English name.
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Vocabulary:
Compound = substances whose contents have fixed ratios and exhibit emergent properties which differ from their elements (two or more)
Compounds are used throughout the world and everywhere, in which their emergent properties are utilized in a variety of ways. For example, salt–a widely-used seasoning, is the combination of sodium and chloride, two elements with individually dangerous properties. Additionally, water is formed through two gases, which, when set into a fixed ratio create this life-giving substance.
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Biologically speaking, compounds usually contain three-to-four different elements, for instance sugar–the body’s pathway to ATP, is made of oxygen, nitrogen, hydrogen, and carbon. The body, because it relies so heavily on these elements through compounds, about 20% of the human body is made up of these organic elements.
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The number of elements needed for life varies from organism to organism. Although plants only need 17, humans require 25 of the natural elements to live. Firstly, oxygen, hydrogen, carbon, and nitrogen, as the building blocks for all molecular proteins, carbohydrates, and lipids, in addition to the calcium and phosphorous needed for bones and teeth, make up about 99% of the body. Subsequently, the remaining 1% goes to the rest of the elements like potassium, sodium, sulfur, and magnesium which used throughout the body in nerve functioning and chemical reactions.
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Vocabulary:
Trace elements = substances found in extremely small amounts
The body has a lot of trace elements comprising the last 0.01% of the human body–iron, zinc, cobalt, silicon, selenium, tin, boron, and chromium are all elements that appear in the body at minute levels. However, this does not diminish their importance at all, as iron is necessary for oxygen transport and therefore life, acting as an essential element for all forms of life. In contrast, other elements are not so ubiquitous, as only certain families, classes, etc needed specific trace elements.
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Some trace elements are difficult to source and access, leading to conditions and effects on the body, such as iodine deficiency. Supplied by foods such as milk, seafood, leafy greens, and kelp, the thyroid gland produces a hormone involving iodine. Therefore when the body does not have enough iodine for the thyroid glands to function properly, goiters or the swelling of thyroid glands occur. Although not usually lethal, iodine deficiency can still lead to many complications especially during pregnancy and childhood, where miscarriages, disabilities and other impacts pose hazards to iodine deficient individuals. In response, efforts in making iodized salt the global standard, have found a modicum of success, however for underdeveloped countries and similar regions, most are unable to access resources. Clearly, the impact of trace elements not only in an individual but also globally can be seen, highlighting the significance of balancing the human body and its nutritional requirements.
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Iron, despite its significance, is another common deficiency usually targeted through diet supplementation and adjustment. Although more prevalent in children and women, undeveloped countries are still affected, with higher numbers of this nutrient disorder as well. Actions like boosting the iron in grocery store bread in the USA are examples of how important adequate nutrition is, and the complications that accompany maintaining such balance.
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A surplus of these trace elements, although not as detrimental to health as deficiencies, still pose palpable threats, for example requiring warnings on iron supplements to ensure that children are not accidentally poisoned from excess iron.
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Other ways to combat deficiencies of these trace elements is by adding them to publicly utilized resources, which was the case for fluorine. All water sources have slightly fluorinated concentrations, which helps to prevent tooth decay and cavities. Here, society and science can be seen collaborating to treat an issue.
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However, despite this, misinformation still creates conflict between the public and science, where suspicion and clashing opinion often overrules the data and evidence supporting movements or scientifically-based actions. For example, fluoridation of all drinking water has sparked controversy as some claim it interferes with their rights or endangers them on account of differing scientific evidence.
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Vocabulary:
Atom = farthest point matter can be broken down into while maintaining its element’s properties (smallest unit of matter)
Atoms, derived from the Greek word for indivisible, comprise matter’s smallest building blocks (for scaling purposes, a million atoms can fit into a printed period). Each element has their own unique atom, these differences producing the properties and characteristics each element is defined by.
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Vocabulary:
subatomic particle = part making up an atom
Proton = atomic component possessing a singular positive charge
Electron = atomic component possessing a singular negative charge
Neutron = atomic component possessing no charge (neutral charge)
Subatomic particles make up the atom, and there are many ways to split atoms in many other types of subatomic particles. However protons, electrons, and neutrons are the most significant; singularly positively charged protons, singularly charged negative electrons, and neutrally charged neutrons are all particles that make up the atom.
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Vocabulary:
Nucleus = center of the atom composed of protons and neutrons
Structurally speaking, the atom consists of two main parts: the nucleus and the negatively charged field of flying electrons orbiting the nucleus. The two most common ways to represent atoms so they can be easily visually understood is through representing the orbiting electrons on a ring around the nucleus, or having the ring turn into a cloud. It is important to remember that these images are made for convenience, as nothing is drawn to scale. Proportionally electrons are much smaller than protons and neutrons, and the ‘cloud’ that the electrons move around is massive compared to the nucleus. (Think baseball stadium to pea to gnat)
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Vocabulary:
Atomic number = amount of protons the atom’s nucleus contains
Atoms and distinguished by their atomic number, which also translates to how many electrons the atom has, as unless specified, any given atoms is assumed to be neutral, therefore having equal positive and negative charges (equal numbers of protons and electrons).
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Vocabulary:
Dalton = measuring unit for the weight of subatomic particles
Mass number = sum of proton and neutron quantity
Atomic mass = weight of the entire atom **electron's mass is extremely low, so mass number and atomic mass are almost equal
The mass number of an atom is also very important, which often correlates to its atomic mass, with both values involving the combined mass of the atom’s protons and neutrons.
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Vocabulary:
Isotope = atoms of the same element with different masses (different amount of neutrons)
An atom’s atomic number defines its elemental identity, however the amount of neutrons does not. This means that the atom’s mass can change and it will remain the same element as before, otherwise known as an isotope. As an example, carbon, with an atomic number of 6, has three well-known isotopes: carbon 12, the most common isotope, carbon 13, much less common, and carbon 14, a very rare isotope. (**note that isotopes can only be named when there are two or more atoms of the same element to compare.)
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Vocabulary:
Radioactive isotope = an isotope whose nucleus is not stable, and will emit random waves of energy as it decays
While some isotopes may be stable, so their nuclei never risk decay, other isotopes are different–radioactive isotopes like carbon 14, as it decays, sends out random waves of its own particles and energy, which poses a danger to nearby molecular organisms. However some radioactive isotopes like carbon 14 can still be helpful, as carbon dating uses carbon 14 to gauge the age of fossils.
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Vocabulary:
Radioactive tracers = certain radioactive isotopes that can be detected using machines to monitor or regulate something
Anything living cannot recognize radioactive isotopes, so any type of isotope can be absorbed by the organism, as long as its elemental identity remains. Scientists and doctors use this property to ‘stain’ people or their subjects to monitor their insides, as the radioactive isotope is easily found by instruments.
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Radioactive tracers also remain intact during chemical reactions, so scientists use these sort of tags to track certain molecules during a chemical reaction. For example, the process of photosynthesis can be monitored if the plant is given carbon 14.
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Within the human body, some chemicals gather in specific regions. Doctors utilize this, and give patients very small amounts of a tracer to create a visualization of that area. Additionally, some radioactive isotopes can be used to treat certain illnesses. For example, radioactive iodine is commonly used to target thyroid cancer, as the radioactive isotopes accumulate in the thyroid gland.
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Vocabulary:
PET = a type of visualizing technology used to detect cancers and heart disorders
Tracers can also be added to food or substances the human body metabolizes. PET can then use those tracers to create images of the body’s interior. This type of technology reveals information of Alzheimer’s as it creates a visual representation of what areas of the brain have been filled with beta-amyloid, a type of protein that fills the brain’s folds.
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Any kind of radioactivity is still very dangerous, and can kill a human by damaging molecular structures such as DNA. The random waves of energy and particles released by unstable atoms threatens lives, especially at higher levels. Some examples of this devastation are in explosions or malfunctions of nuclear reactors, where people with high exposure often died and many had to be evacuated from the area. (Ukraine and Japan)
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Natural substances also pose threats to public health, as radon, a type of gas released by uranium, is also radioactive and therefore exerts the same influence. People can protect themselves by checking their home’s radon levels.
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Vocabulary:
Electron shells = levels on which electrons exist at unique numbers and distances from the nucleus
During chemical reactions, the interactions of atoms only involve electrons and their movement. Electrons are not located within the nucleus, and instead orbit the center, organizing into electrons shells that can be detailed with a diagram. The amount of shells an atom needs depends on its atomic number (The first level can only hold two electrons, the second level can hold 8, etc)
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Vocabulary:
Orbitals = specific areas within in an electron shell that hold up to two electrons
The periodic table lists each element in rows according to how many electron shells they have. The first shell has only one orbital, therefore making its electron capacity two. The second ring has four orbitals, which means it can hold eight electrons.
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Vocabulary:
Valence shell = the outermost electron shell of an atom
Atoms interact with each other depending on how filled their valence shell is; the movement of electrons resulting from the interactions of reactive atoms allow these atoms to fill their valence shells.
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An atom's reactivity is determined by how filled their valence shell is, or how much potential they have of causing the movement of electrons. Some are more reactive than others–inert atoms already have filled valence shells.
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Vocabulary:
Chemical bonds = attractions of atoms towards each other produced by the movement of electrons (caused by the atoms’ incomplete valence shells)
Ionic bonds = attractions created when electrons are transferred between atoms (resulting charge different attracts both atoms towards each other)
Atoms react with one another when there are unfilled valence shells, which then causes the ingress, egress, or sharing of electrons to occur.
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Vocabulary:
Covalent bond = bond formed through sharing a pair of electrons between atoms (more than one pairs can be shared at a time)
Valance = an atom’s potential reactivity/electron movement
Covalent bonds form not through charge differences but the physical sharing of electrons. The amount of covalent bonds that can form between atoms depends on how many electrons an atom’s valence electron shell needs.
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Vocabulary:
Molecule = two or more atoms bonded together through covalent bonds
Hydrogen molecules are made of two hydrogen atoms covalently bonded together, as they both share each other’s singular electron.
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Vocabulary:
Electronegativity = a term that describes how much an atom attracts covalently shared electrons
Nonpolar covalent bond = a covalent bond that involves atoms of the same or close electronegativity, allowing them to share the electrons equally
Other atoms, over the atoms they are covalently bound, may attract the shared electrons more than the other. Electronegativity measures this attraction. In contrast, in the case of a hydrogen molecule, as they are the same type of atom with the same electronegativity, the electrons are shared equally. Bonds like these are called nonpolar covalent bonds.
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Vocabulary:
Polar covalent bond = atoms whose electronegativity differs, so that their covalent bonds involve uneven sharing of electrons
When the electronegativity differs this means that the electrons within a covalent bond are unevenly shared, depending on which atom is more electronegative.
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Molecular formula = shows how many of each atom is present in the molecule (by subscript) and the type of atom (through chemical symbol)
Electron distribution diagram = depicts the way that each atom involved in a chemical bond completes their valence shell
Structural formula = **only for covalent bonds–shown as lines and visualizes model similarly to actual molecule structure
Space-filling model = color-coded balls are the most accurate way to represent individual atoms within a molecule
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Ionic bond = attraction of oppositely charged ions
Ionic bonds occur when the electronegativity of an atom is so much stronger than other, that it actually causes the electron to be pulled away from the less electronegative atom. Thus, the less electronegative atom now becomes an ion with a positive charge, as losing that electron now makes the amount of protons greater, and the more electronegative atom becomes an ion with a negative charge through gaining that additional electron.
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Ion = atom or molecule who gained or lost electrons, resulting in an electrical charge
Table salt is an example of an ionic bond: sodium has one valence electron, and chloride has seven–chlorine is way more electronegative compared to sodium (chlorine needs only one electron; sodium needs seven). As a result chlorine pulls sodium’s singular valence electron away from it, completing its octet on its third shell as an ion with a net electrical charge of 1-, and leaving sodium as an ion with a net electrical charge of 1+ and two electron shells. An attraction is then formed, creating the compound table salt.
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Salt = ionic compounds
Salt crystals can grow to any size with any amount of ions (1:1 ratio of sodium to chloride), but the ratios vary between the types of salt. Environment affects the strength of a salt’s bonds–when dry the bonds are very strong, but when exposed to water salt is dissolved very easily. (Most drugs are a type of salt due to their stability when dry and solubility when wet)
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Hydrogen bond = weak attraction of hydrogen atom with O, N, or F atom
Polar molecule = molecule has uneven charge distribution
Hydrogen bonds are prevalent throughout life, and are very important because of their frequency and weakness. Within biology, they connect the double strands of DNA, help with protein folding, the transfer of information, and also construction of the proteins themselves. They are formed between polar molecules, because a slightly positive hydrogen atom of a molecule will be attracted to a slightly negative O, N, or F atom of a molecule. Water best illustrates hydrogen bonding because it has polar covalent bonds, which gives oxygen a more negative charge and hydrogen a more positive charge. As a result, positive hydrogens (2 max) are attracted to the side with the negative O, and negative oxygens are attracted to the side with positive hydrogens (max of four hydrogen bonds per water molecule).
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Chemical reaction = the assembly and destruction of existing CHEMICAL BONDS
Reactants = the starting materials of a chemical reaction
Product = final substance of a chemical reaction
Chemical reactions are constantly occurring in cells. For example, oxygen and hydrogen react to form water (arrow represents conversion) through covalent bonds. The amount of matter in both the reactants and products are conserved. Matter is not created or destroyed during chemical reactions, they only rearrange the molecular structure of substances/compounds.
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Photosynthesis = chemical process that takes CO2 and H2O and converts it to C6H12O6 and O2 (6 CO2 + 6 H2O → C6H12O6 + 6 O2)
CHemical reactions are vital to life, for example photosynthesis powers all life on earth, as it converts the energy of sunlight with carbon dioxide and water into glucose and oxygen. The amount of matter and energy, as seen, is kept equal on both sides. They are always conserved. Most chemical reactions occur within the primarily water-based interior of a cell.
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Cohesion = sticking of the same kinds of molecules together
Adhesion = sticking of different kinds of molecules together
Capillary action = upwards movement of water through adhesion and cohesion (pulling force from water evaporation)
The most important property of water is its ability to form hydrogen bonds with itself as a result of polar covalent bonds. As a result, water is very cohesive, and is continually breaking and forming new hydrogen bonds so that in every couple trillionths of a second, water molecule hydrogen bonds break apart and are at the same time still connected to other molecules. Adhesion is another important property of water, as it allows water to bond with other molecules. For example, when water evaporates in leaves, it sends a chain reaction of forces that pull more water all the way from the root. The tiny veins in the plants support this capillary action, as their thin cell walls allow the water to more easily adhere to its surface.
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Surface tension = measure the effort needed to stretch or break a liquid’s surface
Water, as a result of its many hydrogen bonds, has very high surface tension This forms a kind of skin on its surface–observable when slightly overfilling a glass of water. Some organisms are able to walk on this very thin film.
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Thermal energy = energy involved in an atom or molecule’s random movement
Heat = movement of energy from warm to cool subject
Temperature = average of thermal energy
A body’s thermal energy pertains to the movement of its particles, which is measured with temperature. Heat describes the transfer of thermal from warm to cool subjects. Some things heat, or atomically/molecularly increase in movement faster than others. For example the metal in a pot will heat much faster compared to the water in it. (Hydrogen bonds affect temperature change)
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Water needs so much more energy to change temperature because of the interference of hydrogen bonds. Energy must be absorbed to break hydrogen bonds, so the water will need a lot of energy to go in and break up those bonds, resulting in only a couple degree changes of temperature. Additionally, energy is released when hydrogen bonds form, so when cooling water it takes much longer on account of the thermal energy and thus heat being released by water molecules as hydrogen bonds re-form.
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Evaporative cooling = decrease in temperature when most excited molecules leave a substance
Because earth has so much water, its climate is relatively well-regulated. The lakes, rivers, and oceans that absorb extra energy from the sun release energy as well during the winter/cooling periods. This creates optimal environments for marine life, as the temperature is so mild. In a similar fashion, water also tempers land climates, as coastal areas usually have more mild climates. Additionally, the human body is 60-70% made of water, which has the same regulating effects. Sweating is an example of this, as when the sweat evaporates, it takes with the thermal energy and therefore heat involved in the evaporation of those molecules, having a cooling effect. This prevents overheating in a lot of organisms.
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The property of ice floating on water is key for the preservation of marine life (top-down formation of ice), and other types of ice-dependent life. Ice acts as an insulator for these organisms. When ice crystals form, the water molecules most assume even and stable hydrogen bonds, which enlarges spaces between the molecules. Compared to liquid water, ice is much more spread apart, making it less dense. This is why it floats in liquid water.
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Solution = liquid with uniform mixture of two or more substances
Solvent = substance that dissolves
Solute = substance being dissolved
Water is also able to dissolve a number of substances, and these water solutions are called aqueous solutions. Water can dissolve any polar substances, as either the slightly positive hydrogens or slightly negative oxygens will be attracted to the molecule. For example, salt is soluble in water because the negatively charged chlorine attracts the hydrogens of water, and the positively sodium attracts oxygens. As a result, the water comes in and separates the Cl and Na from each other as hydrogen bonds are formed, separating the ions. Water is very important for biological functions, as it can dissolve sugars, proteins, etc in a cell’s external and internal environment.
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Acid = substance that increases H+ ion concentration of a solution
Sometimes water molecules dissociate with each other, causing it to separate into H+ and OH-. The presence of this ion and molecule can have very influential impacts on protein function. When something adds H+ ions to a solution, it is considered an acid (hydrochloric acid)
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Base = substance that decreases H+ ion concentration (therefore increases OH- concentration)
Some kinds of basic substances release OH- acceptors into the solution so that it will combine with the H+ and form H20. However other types of alkaline substances will simply accept H+ ions themselves, subsequently increasing the concentration of OH- relative to what had been there before.
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pH = scale that depicting acidic versus alkaline concentrations of a particular substance
Buffer = chemicals that reduce effect of pH changes
0 is the most acidic part of the pH scale, and 14 the most basic. Neutral solutions are at a pH of 7. Some acidic foods are lemons, citruses, etc. The pH of blood must remain at 7.4 if not very close, otherwise the body will not be able to function. This is why buffers are so important, because humans cannot survive more than a couple minutes will an out-of-balance of pH. Buffers work by adjusting the concentration of H+ ions in a solution by either accepting donating more ions into the blood.
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Ocean acidification = decrease of pH in ocean from additional dissolved CO2 (linked to increasing global climate)
Calcification = the animals in coral form calcium carbonate skeletons t=from calcium and carbonate ions
High levels of CO2 in the air also means that the amount of CO2 in the ocean has increased as well. As oceans absorb what thermal energy they are able to, this excess of carbon means that a process involving the solubility of CO2 lowers the pH, making the ocean more and more acidic. The presence of additional CO2 uses up the calcium ions, making bicarbonate ions and lessening the skeleton-making resources available to these coral animals and other animals that rely on constructing shells. It is predicted that this will result in a carbonate ion decrease by 40% by 2100.
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This effect of CO2 on calcification has been observed and experimented by scientists, who have been able to determine that a decrease in carbonate ions will harm coral reefs. A study in a coral reef system in Arizona was able to simulate an environment in which the amount of calcium ions, pH, and temperature were regulated and the concentration of carbonate ions was adjusted. It revealed that the rate of calcification decreased and therefore the coral animals took longer to grow. Some other natural environments reveal how the presence of too much CO2 affects structural complexity and the growth rate of corals. This also affects the reef community, as it is harder to survive when the coral structures are diminished. These consequences of increasing CO2 concentration (resulting in carbonate ion decrease) demonstrate the danger coral reef communities are in from.
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Water is vital for life, and so when trying to discover life on other planets, scientists first look for signs of water. Mars has shown promising signs–polar ice caps, water-rich soil, water-saturated molecules, seasonal streams, and other signs have been collected by the various rovers sent to Mars. This has sparked interest and investigation of the potential for life on other planets.