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Flip Charts

Atoms and Chemical Bonding Flip Chart Set

Physical Science - Middle School

 
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\|xiFFIFGy00531qzZ Copyright © NewPath Learning. All rights reserved. www.newpathlearning.com 34-6922 Charts Charts ATOMS & CHEMICAL BONDING ATOMS & CHEMICAL BONDING Sturdy, Free-Standing Design, Perfect for Learning Centers! Reverse Side Features Questions, Labeling Exercises, Vocabulary Review & more!
Phone: 800-507-0966 Fax: 800-507-0967 www.newpathlearning.com NewPath Learning® products are developed by teachers using research-based principles and are classroom tested. The company’s product line consists of an array of proprietary curriculum review games, workbooks, posters and other print materials. All products are supplemented with web-based activities, assessments and content to provide an engaging means of educating students on key, curriculum-based topics correlated to applicable state and national education standards. Copyright © 2014 NewPath Learning. All Rights Reserved. Printed in the United States of America. Curriculum Mastery® and NewPath Learning® are registered trademarks of NewPath Learning LLC. Science Curriculum Mastery® Flip Charts provide comprehensive coverage of key standards-based curriculum in an illustrated format that is visually appealing, engaging and easy to use. Curriculum Mastery® Flip Charts can be used with the entire classroom, with small groups or by students working independently. Each Curriculum Mastery® Flip Chart Set features 10 double-sided laminated charts covering grade-level specific curriculum content on one side plus write-on/wipe-off charts on reverse side for student use or for small-group instruction. Built-in sturdy free-standing easel for easy display Spiral bound for ease of use Student Activity Guide Ideal for Learning centers In class instruction for interactive presentations and demonstrations Hands-on student use Stand alone reference for review of key science concepts Teaching resource to supplement any program HOW TO USE Classroom Use Each Curriculum Mastery® Flip Chart can be used to graphically introduce or review a topic of interest. Side 1 of each Flip Chart provides graphical representation of key concepts in a concise, grade appropriate reading level for instructing students. The reverse Side 2 of each Flip Chart allows teachers or students to summarize key concepts and assess their understanding. Note: Be sure to use an appropriate dry-erase marker and to test it on a small section of the chart prior to using it. The Activity Guide included provides a black-line master of each Flip Chart which students can use to fill in before, during, or after instruction. While the activities in the guide can be used in conjunction with the Flip Charts, they can also be used individually for review or as a form of assessment or in conjunction with any other related assignment. Learning Centers Each Flip Chart provides students with a quick illustrated view of science curriculum concepts. Students may use these Flip Charts in small group settings along with the corresponding activity pages contained in the guide to learn or review concepts already covered in class. Students may also use these charts as reference while playing the NewPath’s Curriculum Mastery® Games. Independent student use Students can use the hands-on Flip Charts to practice and learn independently by first studying Side 1 of the chart and then using Side 2 of the chart or the corresponding graphical activities contained in the Activity Guide. Reference/Teaching resource Curriculum Mastery® Charts are a great visual supplement to any curriculum or they can be used in conjunction with NewPath’s Curriculum Mastery® Games. Chart # 1: Chart # 2: Chart # 3: Chart # 4: Chart # 5: Chart # 6: Chart # 7: Chart # 8: Chart # 9: Chart #10: Models of the Atom Atomic Configuration & Bonding Chemical Bonding Ionic Bonding Ionic Compounds Covalent Bonding Covalent Compounds Naming Covalent Compounds Metallic Bonding Vocabulary
Models of the Atom © Copyright NewPath Learning. All Rights Reserved. 94-4836 Visit www.newpathlearning.com for Online Learning Resources. Atomic Models Scientists use models to explain things that cannot be seen directly. The first model of the atom was proposed by the Greek philosopher, Democritus, who hypothesized that all matter was made of small particles, called atoms. Dalton’s Atomic Theory In the early 1800s, John Dalton conducted experiments on gases and hypothesized the following: 1) All matter is made of atoms and atoms are small particles that cannot be created, divided or destroyed. 2) All atoms of the same element are identical & different elements have different types of atoms. 3) Atoms combine with other atoms to form new substances. Thomson’s Model of the Atom In 1904, J.J. Thomson proposed a model of the atom as a solid sphere with equal numbers of positive and negative charges spread throughout, much as raisins might be embedded in the surface of a pudding. + Oxygen J.J. Thomson 1856-1940 negative particles - “corpuscles” are repelled by the magnet + + + + + + Thomson’s atomic model “plum pudding” equally distributed positive & negative particles “plum pudding” model Rutherford’s Model In 1911, Ernest Rutherford proposed a new model of the atom which consisted of a positively charged nucleus, containing most of the atomic mass of the atom, and negatively charged electrons orbiting around the nucleus like planets around the Sun. + highest energy level Niels Bohr 1885-1962 lowest energy level photon electron Quantum Mechanical model Erwin Schrödinger 1887-1961 dense areas - more likely to nd electrons probability plot of electron density + Bohr’s Model of the Atom In 1913, Niels Bohr proposed electrons move in stable, or stationary, orbits at fixed distances from the nucleus. Each orbit had an energy associated with it. The closest orbit had the lowest energy and the energy increased with the distance from the nucleus. If an electron moves between orbits, then energy in the form of light (or photons), is absorbed or emitted. Oxygen Ernest Rutherford 1871-1937 gold foil detecting screen alpha-particle emier alpha particles (+ charge) + + + + + + + + + + + + + + + + + + + + + + + + + + + Rutherford’s model atoms in gold foil most alpha particles pass through beam of alpha (+) particles + + + + + + proton quarks nucleus u u d The Modern Atomic Theory In 1926, Erwin Schrödinger used mathematical equations to describe the likelihood of finding an electron in a certain position. This atomic model is known as the Quantum Mechanical model of the atom. This model can be represented as a nucleus surrounded by an electron cloud. The probability of finding the electron is greatest where the cloud is most dense and least likely in a less dense area of the cloud. In 1932, James Chadwick discovered the neutron, an uncharged particle contained in the nucleus. Through continued experimentation, more particles have been discovered in the atom. Quarks are believed to be even smaller units than protons and neutrons. In turn, quarks are made up of vibrating strings of energy. The search to find even smaller particles that make up an atom continues. John Dalton J.J. Thompson Niels Bohr Erwin Schrödinger nucleus
Pause and Review Match the scientist with their atomic model. Fill in the name of the scientist under each model. + + + + + + + + + + + + + + + + Electrons move in stable orbits around the nucleus. Electrons move in an electron cloud around the nucleus. made of called Atoms have a positive nucleus surrounded by negative electrons. The nucleus contains neutral particles called neutrons. Atoms have negatively & positively char ged particles, evenly spread throughout. J.J. Thomson 1856-1940 Niels Bohr 1885-1962 Erwin Schrödinger 1887-1961 James Chadwick 1891-1974 Ernest Rutherford 1871-1937 John Dalton 1766-1844 + + + + + + + + + Electrons move in stable orbits around the nucleus. Electrons move in an electron cloud around the nucleus. Atoms have a positive nucleus surrounded by negative electrons. The nucleus contains neutral particles called neutrons. Niels Bohr 1885-1962 Erwin Schrödinger 1887-1961 James Chadwick 1891-1974 Ernest Rutherford 1871-1937 + + + + + + + in Electrons move in an electron cloud around the nucleus. The nucleus contains neutral particles called neutrons. Schrödinger 1887-1961 James Chadwick 1891-1974 + + + + + + + + + + + + + + + + Electrons move in stable orbits around the nucleus. Electrons move in an electron cloud around the nucleus. All matter is made of small particles called atoms. Atoms have a positive nucleus surrounded by negative electrons. The nucleus contains neutral particles called neutrons. Atoms have negatively & positively char ged particles, evenly spread throughout. J.J. Thomson 1856-1940 Niels Bohr 1885-1962 Erwin Schrödinger 1887-1961 James Chadwick 1891-1974 Ernest Rutherford 1871-1937 John Dalton 1766-1844 + + + + + + + + + + + + + + + + Electrons move in stable orbits around the nucleus. Electrons move in an electron cloud around the nucleus. All matter is made of small particles called atoms. Atoms have a positive nucleus surrounded by negative electrons. The nucleus contains neutral particles called neutrons. Atoms have negatively & positively char ged particles, evenly spread throughout. J.J. Thomson 1856-1940 Niels Bohr 1885-1962 Erwin Schrödinger 1887-1961 James Chadwick 1891-1974 Ernest Rutherford 1871-1937 John Dalton 1766-1844 + + + + + + + + + + + + + + + + Electrons move in stable orbits around the nucleus. Electrons move in an electron cloud around the nucleus. is made of particles called atoms. Atoms have a positive nucleus surrounded by negative electrons. The nucleus contains neutral particles called neutrons. Atoms have negatively & positively char ged particles, evenly spread throughout. J.J. Thomson 1856-1940 Niels Bohr 1885-1962 Erwin Schrödinger 1887-1961 James Chadwick 1891-1974 Ernest Rutherford 1871-1937 John Dalton 1766-1844 + + + + + + + + + Electrons move in stable orbits around the nucleus. Electrons move in an electron cloud around the nucleus. have a positive surrounded negative electrons. The nucleus contains neutral particles called neutrons. Niels Bohr 1885-1962 Erwin Schrödinger 1887-1961 James Chadwick 1891-1974 Ernest Rutherford 1871-1937 Models of the Atom © Copyright NewPath Learning. All Rights Reserved. 94-4836 Visit www.newpathlearning.com for Online Learning Resources.
Atomic Configuration & Bonding © Copyright NewPath Learning. All Rights Reserved. 94-4837 Visit www.newpathlearning.com for Online Learning Resources. Electron Configuration The specific way the electrons are arranged in an atom is called the electron configuration. Electrons play an important role in how elements interact with each other and form compounds. The electrons are distributed among orbital shells or energy levels (1, 2, 3 and so on) that are different distances from the nucleus. The larger the number of the energy level, the farther it is from the nucleus. Electrons that are in the highest or outmost energy level are called valence electrons. The valence electrons are the ones that are lost, gained or shared during chemical bonding. Oxygen atom (neutral) electrons orbital shell 1 orbital shell 2 atomic number = 8 protons = 8 electrons = 8 electron cloud nucleus valence electron valence shell Br 35 Bromine Cl 17 Chlorine F 9 Fluorine I 53 Iodine At 85 Astatine Li 3 Lithium Na 11 Sodium K 19 Potassium Rb 37 Rubidium Cs 55 Cesium Alkali metals 1 valence electron lose electron Halogens 7 valence electrons gain electron Sodium Chlorine Aluminum Oxygen Atoms with 4 or more valence electrons gain electrons. Atoms with 3 or less valence electrons lose electrons. Atoms with full valence shells will not combine with other elements. Neon Electron Dot Diagram The number of valence electrons in an atom of an element determines many properties of that element, including the ways in which the atom can bond with other atoms. An electron dot diagram is often used to depict the valence electrons in an atom. Each atom of an element has a specific number of valence electrons, ranging from 1 to 8. The electron dot diagram includes the element symbol, surrounded by dots. Each dot represents 1 valence electron. The dots are spaced out above, below, to the left, and to the right of the symbol for the first 4 valence electrons. For atoms with greater than 4 valence electrons, the dots must be paired up. The dot diagrams for atoms can also be used to show the bond between different atoms in a molecule. C O Ar valence electrons Carbon Oxygen Argon He 2 Helium Ne 10 Neon Ar 18 Argon Kr 36 Krypton Xe 54 Xenon Rn 86 Radon Argon Neon Krypton Noble gases 8 valence electrons Stability of Atoms Atoms of most elements are more stable and are less likely to react with other atoms, when they have 8 valence electrons in their outer shell. For example, atoms of neon, argon, krypton, and xenon are very unreactive because they all have 8 valence electrons. Atoms usually react in a way that makes each atom more stable by losing, gaining, or sharing electrons in a chemical bond with other atoms. Valence Electrons & Bonding The outermost orbital shell, called the valence shell, is most often involved in chemical bonding. Elements in the same group in the periodic table have the same number of electrons in their valence shell. For example, all elements in group 1, alkali metals, have 1 valence electron. Group 1 atoms prefer to lose 1 electron to become stable. However, all elements in group 17, halogens, have 7 valence electrons. These atoms will gain 1 electron to fill their valence shell.
Pause and Review Fill in the table below. Group # C Cu Be O Xe # of electrons # of valence electrons Element (neutral) Use the information in the illustrations to answer the questions. Select the best answer(s) and then click on the check button. Please make the buttons as large as possible so they can count electrons Na F Ne O 1) Which element has 1 valence electron? ________________________________ 2) Which element is most reactive? ______________________________________ 3) Which element has a full valence shell? ________________________________ 4) Which element is the most stable? ____________________________________ 5) Which element has 6 valence electrons? _______________________________ 6) Which element has 7 valence electrons? _______________________________ Atomic Configuration & Bonding © Copyright NewPath Learning. All Rights Reserved. 94-4837 Visit www.newpathlearning.com for Online Learning Resources.
metal nonmetal Chemical Bonding © Copyright NewPath Learning. All Rights Reserved. 94-4838 Visit www.newpathlearning.com for Online Learning Resources. Chemical Bonds In 1916, Gilbert Newton Lewis, an American scientist, proposed that chemical bonds are formed due to the electron interaction between atoms. His work established the basis of what we know today about chemical bonding. Atoms combine with other atoms through chemical bonds, which result from the strong attractive forces that exist between the atoms. Atoms bond together to become more stable by having a full valence shell. Certain elements are more reactive than others and will be more likely to bond with other elements. There are three main types of chemical bonding: covalent bonding, ionic bonding and metallic bonding. Covalent Bonding The bond between two nonmetals is usually a covalent bond. By sharing electrons, two atoms can mutually complete their valence shells to become more stable. Ionic Bonding The bond between a metal and nonmetal atom is an ionic bond. In this example, to become stable, the metal sodium atom loses one electron in its outer shell, which is gained by the nonmetal chlorine atom, which also becomes stable with a full outer shell electron configuration. H Cl HCl shared electrons Cl- Na+ Cl Na + Cl- Na+ Cl Na + Metallic Bonding Metal atoms bond by forming a metallic bond. The valence electrons in a metallic substance continually move throughout the metal from one atom to another. The atoms that the electrons leave behind become positive ions. The interaction between such ions and valence electrons provides the bonding force that holds a metallic structure together. + + + + + + + + + + + + Al+ ion loose valence electrons loose valence electrons shared electrons
Chemical Bonding Pause and Review Describe and illustrate the three types of chemical bonding. © Copyright NewPath Learning. All Rights Reserved. 94-4838 Visit www.newpathlearning.com for Online Learning Resources. 1. Covalent Bonding ___________________________________ ___________________________________ ___________________________________ ___________________________________ ___________________________________ ___________________________________ 2. Ionic Bonding ___________________________________ ___________________________________ ___________________________________ ___________________________________ ___________________________________ ___________________________________ 3. Metallic Bonding ___________________________________ ___________________________________ ___________________________________ ___________________________________ ___________________________________ ___________________________________
Ionic Bonding © Copyright NewPath Learning. All Rights Reserved. 94-4839 Visit www.newpathlearning.com for Online Learning Resources. Charged Particles A normal atom has a neutral charge with an equal number of protons and electrons. For example, a neutral atom of sodium has eleven protons and eleven electrons. However, since atoms are more stable with a full valence shell, an atom can gain or lose electrons to become stable. If a sodium atom loses 1 electron, it will have a full valence shell but it will no longer be neutral. The atom will now have 11 protons and 10 electrons for a net charge of +1. A neutral atom becomes charged when there are an unequal number of protons and electrons. For example, calcium loses two valence electrons, to form a positive ion with a +2 net charge which is expressed as Ca2+. Alternatively, an oxygen atom gains two electrons to form a negative ion with a -2 charge, which is expressed as O2-. Calcium Oxygen 8 protons 10 electrons -2 net charge O2- 20 protons 18 electrons +2 net charge Ca2+ Ionic Bonds Ionic bonds form as a result of the attraction between positive and negative ions. Ionic bonding normally occurs between metal atoms and nonmetal atoms. Atoms with partially filled outer shells are unstable. To become stable, the metal atom loses one or more electrons in its outer shell, forming a positively charged ion or cation with a stable electron configuration. These electrons are then gained by the nonmetal atom, causing it to form a negatively charged ion or anion, also with a stable electron configuration. The attraction between the oppositely charged ions causes them to come together and form an ionic bond. In this example, sodium (Na) and chlorine (Cl) ions are attracted to each other in a 1:1 ratio and combine to form sodium chloride (NaCl), common table salt. creates a full valence shell Sodium ion (+) metal Chlorine ion (-) non-metal (cation) (anion) Cl Na + forms an ionic bond table salt Cl- Na+ Forming Ions A neutral atom that becomes charged by gaining or losing electrons is called an ion. An atom can acquire a positive or negative charge depending on whether the number of electrons is greater or less than the number of protons in the atom. An atom with more electrons than protons, has a negative charge and is a negative ion, or anion. However, an atom with more protons than electrons, has a net positive charge and is a positive ion, or a cation. Since an electron and a proton have equal but opposite unit charges, the charge of an ion is always expressed as a whole number of unit charges and is either positive or negative. Sodium atom (neutral) single valence electron 11 protons 11 electrons 11+ (-11) = 0 0 neutral charge full valence shell Sodium ion (+1) 11 protons 10 electrons 11+ (-10) = +1 +1 net charge
Ionic Bonding © Copyright NewPath Learning. All Rights Reserved. 94-4839 Visit www.newpathlearning.com for Online Learning Resources. Pause and Review Describe and illustrate ionic bonding. _____________________________________ _____________________________________ _____________________________________ _____________________________________ _____________________________________ Describe and illustrate the formation of ions. _____________________________________ _____________________________________ _____________________________________ _____________________________________ _____________________________________ Describe and illustrate a charged particle. _____________________________________ _____________________________________ _____________________________________ _____________________________________ _____________________________________
Ionic Compounds © Copyright NewPath Learning. All Rights Reserved. 94-4840 Visit www.newpathlearning.com for Online Learning Resources. Naming Ionic Compounds The names of ionic compounds are written by listing the name of the positive ion followed by the name of the negative ion. Therefore, a series of rules is needed to name the positive and negative ions before we can name these compounds. Sodium Chloride Potassium Chloride Aluminum Phosphide Magnesium Sulde Cl Na + P Al 3+ 3– Cl K + S Mg 2+ 2– 1+ 1– += 0 2+ 2– += 0 1+ 1– += 0 3+ 3– += 0 Sodium Chloride Potassium Chloride Aluminum Phosphide Magnesium Sulde Cl Na + P Al 3+ 3– Cl K + S Mg 2+ 2– 1+ 1– += 0 2+ 2– += 0 1+ 1– += 0 3+ 3– += 0 Single atomic positive ions have the name of the element from which they are formed. However, since some metals form positive ions in more than one oxidation state, (ex. Fe2+, Fe3+) the charge on the ion is indicated by a Roman numeral in parentheses immediately after the name of the element (Ex. iron(II), iron(III)). single atomic negative ions Oxide Fluoride Chloride O2– F– Cl– Sul de S2– Bromide Br– Iodide I– Nitrate Sulfate SO 4 2– Chlorate ClO 3 NO 3 Nitrite Sul te SO 3 2– Chlorite ClO 2 NO 2 polyatomic negative ions low oxidation state high oxidation state Hydrogen Sodium single atomic positive ions Potassium H+ Na+ K+ Calcium Ca2+ Iron (II) Fe2+ Iron (III) Fe3+ polyatomic positive ions Hydronium H 3O + Ammonium NH 4 + Hydrogen Sodium single atomic positive ions Potassium H+ Na+ K+ Calcium Ca2+ Iron (II) Fe2+ Iron (III) Fe3+ polyatomic positive ions Hydronium H 3O + Ammonium NH 4 + Negative ions that consist of a single atom are named by adding the suffix “-ide” to the stem of the name of the element. Polyatomic Negative Ions The name of polyatomic negative ions usually ends in either “-ite” or “- ate.” The “-ite” ending indicates a low oxidation state (nitrite ion NO2-). The “–ate” ending indicates a high oxidation state (nitrate ion NO3-). Oxidation State Oxidation state shows the total number of electrons that an atom has gained or lost in order to form a chemical bond with another atom. Writing Formulas for Ionic Compounds When an ionic compound is formed, the ions must combine in a way that the total charges of the compound equal zero. For example, let’s write the correct formula for magnesium chloride. The first step is to write the formulas for the cation Mg2+ and anion Cl-. Next, drop the positive and negative signs; crisscross the superscripts so that they become subscripts and reduce when possible by finding the least common multiple. In this example, the two chlorine ions, with a total charge of -2, balance the +2 charge of the magnesium ion. The cation is always listed first before the anion, resulting in the formula MgCl 2. +3 2(-2) +5 3(-2) oxidation state low oxidation state high oxidation state Nitrate NO 3 Nitrite NO 2 NO 2 NO 3 N +3 +3 2(-2) +5 3(-2) oxidation state low oxidation state high oxidation state Nitrate NO 3 Nitrite NO 2 NO 2 NO 3 N +3 Mg2 Mg2+ Cl– Magnesium Chloride anion (-) cation (+) 1 2 Cl 2 Mg 2 Cl 3 2+ 2(–) += 0 Polyatomic (more than two atoms) positive ions often have common names ending with the suffix “-onium” such as hydr onium (H3O+) or ammonium (NH4+). single atomic negative ions Oxide Fluoride Chloride O2– F– Cl– Sul de S2– Bromide Br– Iodide I– Nitrate Sulfate SO 4 2– Chlorate ClO 3 NO 3 Nitrite Sul te SO 3 2– Chlorite ClO 2 NO 2 polyatomic negative ions low oxidation state high oxidation state
Ionic Compounds Pause and Review Name the following compounds. 1) K2S __________________________________________________________________ 2) CaBr2 _________________________________________________________________ 3) HI ____________________________________________________________________ 4) SrF2 __________________________________________________________________ 5) Be3P2 _________________________________________________________________ 6) Al2O3 _________________________________________________________________ 7) BCl3 __________________________________________________________________ Write the formula for the following ionic compounds. 1) Magnesium Chloride ___________________________________________________ 2) Sodium Nitride ________________________________________________________ 3) Sodium Chlorate _______________________________________________________ 4) Potassium Nitrate ______________________________________________________ 5) Calcium Chlorate ______________________________________________________ 6) Aluminum Nitrite ______________________________________________________ 7) Aluminum Nitrate _____________________________________________________ © Copyright NewPath Learning. All Rights Reserved. 94-4840 Visit www.newpathlearning.com for Online Learning Resources.
Chlorine Cl 2 complete outer shells Covalent Bonding © Copyright NewPath Learning. All Rights Reserved. 94-4841 Visit www.newpathlearning.com for Online Learning Resources. Covalent Bonds Covalent bonding is a type of chemical bonding that occurs between two non-metallic atoms. It is characterized by the sharing of one or more pairs of electrons between atoms. By sharing electrons, two atoms can mutually complete their valence shells to become more stable. For example, since each chlorine atom has 7 electrons in its outer shell, two chlorine atoms will each share an electron to obtain a complete outer shell and form a stable Cl2 molecule. For every pair of electrons shared between two atoms, a single covalent bond is formed. C O O NN double covalent bond triple covalent bond Carbon Dioxide CO 2 Nitrogen gas N 2 C O O NN NN C O O C O O NN double covalent bond triple covalent bond Carbon Dioxide CO 2 Nitrogen gas N 2 C O O NN NN C O O Atoms can also share two or three pairs of electrons and are named accordingly as double and triple bonds. A single line indicates a bond between two atoms, double lines (=) indicate a double bond and triple lines ( ) represent a triple bond. Nonpolar & Polar Co valent Bonding There are two types of covalent bonding - nonpolar and polar. Nonpolar bonding results when two identical non- metals equally share electrons between them. Diatomic molecules such as O 2 or I 2 form nonpolar covalent bonds where both atoms share the electrons equally. Certain other compounds, such as ethane (C2H6), have both polar and nonpolar bonds. Ethane, has polar bonds between the carbon and hydrogen, and nonpolar bonds between the two carbon atoms. C O O + –– polar covalent CO 2 double bond C O O Carbon Dioxide H H H N + + + NH 3 polar covalent Ammonia N H H H O H H O H + + Hydrogen Oxygen H 2O - water molecule polar covalent H H H O O H H H H H H C C –– + + + + + + Ethane - C 2H6 polar covalent bonds nonpolar bond C C H H H H H H Nonpolar Covalent Bond Polar Covalent Bonds CO 2 NH 3 H 2O Ethane - C 2H6 Polar bonding results when two different non-metals unequally share electrons between them. Compounds such as carbon dioxide, ammonia, and water have polar covalent bonds. double bond
Pause and Review Complete the graphic organizers below. Covalent Bonding © Copyright NewPath Learning. All Rights Reserved. 94-4841 Visit www.newpathlearning.com for Online Learning Resources. Covalent Bonding Types of Covalent Bonds include to become results when results when examples include examples include which results in that occurs by is a type of
Covalent Compounds © Copyright NewPath Learning. All Rights Reserved. 94-4842 Visit www.newpathlearning.com for Online Learning Resources. Covalent Compounds The atoms in covalent compounds, also known as molecular compounds, are bonded together by covalent bonds. Unlike ionic compounds which form a regular pattern, covalent compounds form individual molecules that are not connected to each other. Due to weak intermolecular forces, most covalent molecules or covalent compounds are liquids or gases at room temperature, with low melting and boiling points. Although most covalent compounds are gases or liquids at room temperature, there’s a class of solid compounds known as covalent network solids that are bonded by covalent bonds, but in a lattice structure. Such compounds are typically hard, transparent, and have high melting points. Examples include diamond, quartz and graphite, among others. Since covalent molecules do not separate into ions when dissolved in water, they are poor conductors of electricity. diamondquartz graphite H 2O O 2 weakintermolecular forces liquidorgasatroom temperature lowmelting&boiling points donotseparateintoions inasolution poorconductorsof electricity diamond quartz graphite
Pause and Review Complete the concept maps below. Covalent Compounds © Copyright NewPath Learning. All Rights Reserved. 94-4842 Visit www.newpathlearning.com for Online Learning Resources. Characteristics of Covalent Compounds Examples of Solid Covalent Compounds
Naming Covalent Compounds © Copyright NewPath Learning. All Rights Reserved. 94-4843 Visit www.newpathlearning.com for Online Learning Resources. Covalent Formulas The molecular formula for a covalent compound can be derived from its name by writing the symbols for the first and second element and translating the prefixes into subscripts. For example, sulfur trioxide would be written as SO 3. Naming Covalent Compounds Covalent compounds are named in a similar manner to binary ionic compounds. (Binary compounds are compounds made up of only two elements). To name binary covalent compounds apply the following rules: Sulfur Trioxide Carbon Dioxide Nitrogen Triiodide Carbon Tetrachloride SO 3 CO 2 NI 3 CCl 4 Sulfur Trioxide Carbon Dioxide Nitrogen Triiodide Carbon Tetrachloride SO 3 CO 2 NI 3 CCl 4 Sulfur Trioxide Carbon Dioxide Nitrogen Triiodide Carbon Tetrachloride SO 3 CO 2 NI 3 CCl 4 Sulfur Trioxide Carbon Dioxide Nitrogen Triiodide Carbon Tetrachloride SO 3 CO 2 NI 3 CCl 4 Naming binary