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ACTIVIDAD 4. Descubre el mapa del agua en tu comunidad Te invitamos a convertirte en un explorador del agua. Ahora que ya aprendiste de dónde viene, cómo la usamos y por qué es tan importante cuidarla, es momento de mirar a tu alrededor y descubrir cómo el agua forma parte de tu vida diaria. Observa con atención tu comunidad y elabora un mapa del agua, donde muestres cómo se mueve el agua que está presente en tu entorno. Materiales sugeridos: • Hojas reutilizadas, cartulina o cuadernos. • Lápices, colores, marcadores o crayones. • Regla, adhesivos o recortes (opcional). ¿Qué debes hacer? En una hoja o cartulina, dibuja un mapa de tu comunidad o del entorno de tu escuela. No tiene que ser perfecto ni exacto, lo más importante es observar, pensar y representar lo que conoces. ¿Qué puedes incluir en tu mapa? • Lugares donde hay agua: Ríos, quebradas, lagunas, canales o el mar (si están cerca). Grifos, estanques, pozos o bebederos. Plantas de tratamiento de agua potable o de aguas residuales (si conoces alguna). • El recorrido del agua: Trata de averiguar de dónde viene el agua que llega a tu casa o escuela, cómo llega hasta ahí y qué pasa con el agua después de que es usada. • Cuidado del agua: Marca con dibujos o símbolos los lugares donde el agua se cuida, también indica los lugares donde podría desperdiciarse o contaminarse y añade ideas o acciones para proteger mejor el agua en tu comunidad. Reflexiona mientras dibujas: ¿De dónde viene el agua que usas cada día? ¿Qué acciones realizamos para no desperdiciarla? ¿Qué podríamos hacer para proteger mejor el agua en nuestra comunidad? Cuando termines tu mapa, compártelo con tus compañeros y cuéntales lo que descubriste sobre el agua. Juntos pueden crear un diario mural en la escuela para compartirlo con la comunidad y promover grandes cambios. 2.2.1 ¿Cómo funciona una planta potabilizadora? Para entender cómo el agua pasa de un río o un embalse hasta el grifo de tu casa siendo totalmente segura, podemos imaginar la planta potabilizadora como una gran fábrica de limpieza que utiliza procesos físicos y químicos, de acuerdo a los siguientes pasos: 1) Captación: El primer paso es extraer el agua de la fuente natural. En la entrada de la planta hay rejas de distintos tamaños que funcionan como un filtro gigante, separando objetos grandes como ramas, plásticos o piedras para evitar que dañen la maquinaria de la planta. 2) Coagulación y floculación: Se añaden sustancias químicas que facilitan la unión de las partículas pequeñas para que luego formen grumos más grandes, llamados flóculos, que son más fáciles de separar del agua. 3) Decantación: Una vez que la suciedad se ha agrupado en flóculos más pesados, el agua pasa a grandes tanques, donde por efecto de la gravedad, esos flóculos se depositan en el fondo y forman un lodo, mientras que el agua más limpia queda en la parte superior y continúa el proceso. 4) Filtración: Aunque el agua ya parezca limpia, aún puede tener impurezas muy pequeñas. Para eliminarlas, el agua atraviesa capas de arena y otros materiales como carbón activado, que actúan como filtros. 5) Desinfección: Este es el paso final para garantizar que el agua no nos enferme, pues se eliminan microorganismos, bacterias y virus, para ello se añade una cantidad controlada de cloro o se utiliza luz ultravioleta (UV) u ozono. 6) Análisis de laboratorio: Se realizan análisis físicos y químicos para asegurar la calidad del agua. Gracias a las plantas potabilizadoras y al trabajo de muchas personas, el agua llega a nuestras casas limpia y segura. Sin embargo, el agua es un recurso limitado. Aunque la tecnología de las plantas es muy avanzada, este proceso requiere mucha energía, conocimientos y cuidado, por lo que proteger y usar el agua de forma responsable es tarea de todos. Aprende más de la potabilización del agua con Veolia: https://www.youtube.com/watch?v=bmtDt2yHwnQ 2.3 Detectives del agua: ¿Qué pasa con el agua después de usarla? Después de usar el agua en casa, por ejemplo, al lavarnos las manos, ducharnos o utilizar el inodoro, el agua no desaparece. Se convierte en agua usada y comienza un nuevo recorrido dentro del ciclo urbano, tal como se mencionó anteriormente

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How to use prepositions of place in Vietnamese| Learn Southern Vietnamese With SVFF - YouTube https://www.youtube.com/watch?v=88FwjYFW8Pc Transcript: (00:02) tiếng Việt giọng miền nam cho người nước ngoài chào mọi người Mình là Thu Trang Today We will learn about Seasons of place [âm nhạc] Chị ơi cho em hỏi em hỏi gì quay bán đồng hồ ở đâu vậy chị khu mua sắm này to quá quầy bán đồng hồ quầy bán đồng hồ ở đối diện quầy bán quần áo ở giữa mấy cái quầy bán nữ trang bên phải là quầy bán bánh ngon lắm nha ở đằng sau á có quầy bán cà phê đã hàng giảm giá mua 2 tặng 1 đó còn ở bên trái thì (01:07) từ đây tới đó xa không chị không xa lắm đâu Chỉ Khoảng vài ba bước chân thôi à Trời đất nhưng em mới hỏi một người này người đó nói đi vòng vòng xe lắc Không vứt phải đâu em đi thẳng đường này á là tới liền cảm ơn chị nhiều nha Không có gì kia 1 bên trái nhà hàng ở bên trái bưu điện số 2 bên phải ngân hàng ở bên phải công viên số 3 ở trước infant (02:10) tiệm cà phê ở trước khách sạn số 4 ở sau bãi đậu xe ở sau công viên số 5 đối diện opec trường học ở đối diện nhà sách số 6 ở bên trong nhà hàng Nhật ở bên trong khu mua sắm số 7 ở bên ngoài có rất nhiều xe hơi đậu ở bên ngoài quán trà sữa số 8 ở giữa mệnh vị ở giữa quán cà phê và bưu điện số 9 kế bên ngân hàng ở kế bên rạp chiếu phim số 10 main (03:20) được oflox the street siêu thị ở bên kia đường [âm nhạc] Where is the Places [âm nhạc] [âm nhạc] Chào bạn [âm nhạc] tao sẽ giật mình (04:43) [âm nhạc]

Understanding Quantum Theory of Electrons in Atoms The goal of this section is to understand the electron orbitals (location of electrons in atoms), their different energies, and other properties. The use of quantum theory provides the best understanding to these topics. This knowledge is a precursor to chemical bonding. As was described previously, electrons in atoms can exist only on discrete energy levels but not between them. It is said that the energy of an electron in an atom is quantized, that is, it can be equal only to certain specific values and can jump from one energy level to another but not transition smoothly or stay between these levels. The energy levels are labeled with an n value, where n = 1, 2, 3, …. Generally speaking, the energy of an electron in an atom is greater for greater values of n. This number, n, is referred to as the principal quantum number. The principal quantum number defines the location of the energy level. It is essentially the same concept as the n in the Bohr atom description. Another name for the principal quantum number is the shell number. The shells of an atom can be thought of concentric circles radiating out from the nucleus. The electrons that belong to a specific shell are most likely to be found within the corresponding circular area. The further we proceed from the nucleus, the higher the shell number, and so the higher the energy level (Figure 9.4.1). The positively charged protons in the nucleus stabilize the electronic orbitals by electrostatic attraction between the positive charges of the protons and the negative charges of the electrons. So the further away the electron is from the nucleus, the greater the energy it has. This quantum mechanical model for where electrons reside in an atom can be used to look at electronic transitions, the events when an electron moves from one energy level to another. If the transition is to a higher energy level, energy is absorbed, and the energy change has a positive value. To obtain the amount of energy necessary for the transition to a higher energy level, a photon is absorbed by the atom. A transition to a lower energy level involves a release of energy, and the energy change is negative. This process is accompanied by emission of a photon by the atom. The following equation summarizes these relationships and is based on the hydrogen atom: The values nf and ni are the final and initial energy states of the electron. The principal quantum number is one of three quantum numbers used to characterize an orbital. An atomic orbital, which is distinct from an orbit, is a general region in an atom within which an electron is most probable to reside. The quantum mechanical model specifies the probability of finding an electron in the three-dimensional space around the nucleus and is based on solutions of the Schrödinger equation. In addition, the principal quantum number defines the energy of an electron in a hydrogen or hydrogen-like atom or an ion (an atom or an ion with only one electron) and the general region in which discrete energy levels of electrons in a multi-electron atoms and ions are located. Another quantum number is l, the angular momentum quantum number. It is an integer that defines the shape of the orbital, and takes on the values, l = 0, 1, 2, …, n – 1. This means that an orbital with n = 1 can have only one value of l, l = 0, whereas n = 2 permits l = 0 and l = 1, and so on. The principal quantum number defines the general size and energy of the orbital. The l value specifies the shape of the orbital. Orbitals with the same value of l form a subshell. In addition, the greater the angular momentum quantum number, the greater is the angular momentum of an electron at this orbital. Orbitals with l = 0 are called s orbitals (or the s subshells). The value l = 1 corresponds to the p orbitals. For a given n, p orbitals constitute a p subshell (e.g., 3p if n = 3). The orbitals with l = 2 are called the d orbitals, followed by the f-, g-, and h-orbitals for l = 3, 4, 5, and there are higher values we will not consider. There are certain distances from the nucleus at which the probability density of finding an electron located at a particular orbital is zero. In other words, the value of the wavefunction ψ is zero at this distance for this orbital. Such a value of radius r is called a radial node. The number of radial nodes in an orbital is n – l – 1. Consider the examples in Figure 9.4.2. The orbitals depicted are of the s type, thus l = 0 for all of them. It can be seen from the graphs of the probability densities that there are 1 – 0 – 1 = 0 places where the density is zero (nodes) for 1s (n = 1), 2 – 0 – 1 = 1 node for 2s, and 3 – 0 – 1 = 2 nodes for the 3s orbitals. The s subshell electron density distribution is spherical and the p subshell has a dumbbell shape. The d and f orbitals are more complex. These shapes represent the three-dimensional regions within which the electron is likely to be found. Principal quantum number (n) & Orbital angular momentum (l): The Orbital Subshell: https://youtu.be/ms7WR149fAY If an electron has an angular momentum (l ≠ 0), then this vector can point in different directions. In addition, the z component of the angular momentum can have more than one value. This means that if a magnetic field is applied in the z direction, orbitals with different values of the z component of the angular momentum will have different energies resulting from interacting with the field. The magnetic quantum number, called ml, specifies the z component of the angular momentum for a particular orbital. For example, for an s orbital, l = 0, and the only value of ml is zero. For p orbitals, l = 1, and ml can be equal to –1, 0, or +1. Generally speaking, ml can be equal to –l, –(l – 1), …, –1, 0, +1, …, (l – 1), l. The total number of possible orbitals with the same value of l (a subshell) is 2l + 1. Thus, there is one s-orbital for ml = 0, there are three p-orbitals for ml = 1, five d-orbitals for ml = 2, seven f-orbitals for ml = 3, and so forth. The principal quantum number defines the general value of the electronic energy. The angular momentum quantum number determines the shape of the orbital. And the magnetic quantum number specifies orientation of the orbital in space, as can be seen in Figure 9.4.3. Figure 9.4.4 illustrates the energy levels for various orbitals. The number before the orbital name (such as 2s, 3p, and so forth) stands for the principal quantum number, n. The letter in the orbital name defines the subshell with a specific angular momentum quantum number l = 0 for s orbitals, 1 for p orbitals, 2 for d orbitals. Finally, there are more than one possible orbitals for l ≥ 1, each corresponding to a specific value of ml. In the case of a hydrogen atom or a one-electron ion (such as He+, Li2+, and so on), energies of all the orbitals with the same n are the same. This is called a degeneracy, and the energy levels for the same principal quantum number, n, are called degenerate energy levels. However, in atoms with more than one electron, this degeneracy is eliminated by the electron–electron interactions, and orbitals that belong to different subshells have different energies. Orbitals within the same subshell (for example ns, np, nd, nf, such as 2p, 3s) are still degenerate and have the same energy. While the three quantum numbers discussed in the previous paragraphs work well for describing electron orbitals, some experiments showed that they were not sufficient to explain all observed results. It was demonstrated in the 1920s that when hydrogen-line spectra are examined at extremely high resolution, some lines are actually not single peaks but, rather, pairs of closely spaced lines. This is the so-called fine structure of the spectrum, and it implies that there are additional small differences in energies of electrons even when they are located in the same orbital. These observations led Samuel Goudsmit and George Uhlenbeck to propose that electrons have a fourth quantum number. They called this the spin quantum number, or ms. The other three quantum numbers, n, l, and ml, are properties of specific atomic orbitals that also define in what part of the space an electron is most likely to be located. Orbitals are a result of solving the Schrödinger equation for electrons in atoms. The electron spin is a different kind of property. It is a completely quantum phenomenon with no analogues in the classical realm. In addition, it cannot be derived from solving the Schrödinger equation and is not related to the normal spatial coordinates (such as the Cartesian x, y, and z). Electron spin describes an intrinsic electron “rotation” or “spinning.” Each electron acts as a tiny magnet or a tiny rotating object with an angular momentum, even though this rotation cannot be observed in terms of the spatial coordinates. The magnitude of the overall electron spin can only have one value, and an electron can only “spin” in one of two quantized states. One is termed the α state, with the z component of the spin being in the positive direction of the z axis. This corresponds to the spin quantum number ms=12. The other is called the β state, with the z component of the spin being negative and ms=−12. Any electron, regardless of the atomic orbital it is located in, can only have one of those two values of the spin quantum number. The energies of electrons having ms=−12 and ms=12 are different if an external magnetic field is applied. Figure 9.4.5 illustrates this phenomenon. An electron acts like a tiny magnet. Its moment is directed up (in the positive direction of the z axis) for the 12 spin quantum number and down (in the negative z direction) for the spin quantum number of −12. A magnet has a lower energy if its magnetic moment is aligned with the external magnetic field (the left electron) and a higher energy for the magnetic moment being opposite to the applied field. This is why an electron with ms=12 has a slightly lower energy in an external field in the positive z direction, and an electron with ms=−12 has a slightly higher energy in the same field. This is true even for an electron occupying the same orbital in an atom. A spectral line corresponding to a transition for electrons from the same orbital but with different spin quantum numbers has two possible values of energy; thus, the line in the spectrum will show a fine structure splitting. The Pauli Exclusion Principle An electron in an atom is completely described by four quantum numbers: n, l, ml, and ms. The first three quantum numbers define the orbital and the fourth quantum number describes the intrinsic electron property called spin. An Austrian physicist Wolfgang Pauli formulated a general principle that gives the last piece of information that we need to understand the general behavior of electrons in atoms. The Pauli exclusion principle can be formulated as follows: No two electrons in the same atom can have exactly the same set of all the four quantum numbers. What this means is that electrons can share the same orbital (the same set of the quantum numbers n, l, and ml), but only if their spin quantum numbers ms have different values. Since the spin quantum number can only have two values (±12), no more than two electrons can occupy the same orbital (and if two electrons are located in the same orbital, they must have opposite spins). Therefore, any atomic orbital can be populated by only zero, one, or two electrons. The properties and meaning of the quantum numbers of electrons in atoms are briefly

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